Deck 11: Solutions: Properties and Behavior

Question 1

Which graph best describes how the vapor pressure of a substance varies with temperature according to the Clausius-Clapeyron equation? ln(P) is plotted on the y-axis, and 1/T is plotted on the x-axis. The origin (0, 0) is not necessarily located where the axes cross.

A) I
B) II
C) III
D) IV

Answer 1

II

Question 2

What is the vapor pressure of an aqueous solution that has a solute mol fraction of = 0.100? The vapor pressure of water is 25.756 mm Hg at 25°C.

A) 23.2 mm Hg
B) 2.58 mm Hg
C) 25.8 mm Hg
D) 0.900 mm Hg
E) 22.3 mm Hg

Answer 2

23.2 mm Hg

Question 3

Which intermolecular interactions are likely to result in the highest vapor pressure of a substance?

A) dispersion interactions
B) dipole-dipole interactions
C) hydrogen bonds
D) ion-dipole interactions
E) ion-ion interactions

Answer 3

dispersion interactions

Question 4

Which statement about the vapor pressure of a liquid is not correct?

A) The vapor pressure increases with increasing temperature because more molecules have sufficient energy to overcome the attractive forces that hold them in the liquid.
B) The vapor pressure decreases with increasing strength of intermolecular interactions.
C) The pressure at which the rate of molecules evaporating equals the rate of molecules condensing is called the vapor pressure.
D) Adding a nonvolatile solute to a liquid increases the vapor pressure.
E) The enthalpy of vaporization determines how sensitive the vapor pressure of a substance is to a change in temperature.

Question 5

Indicate which aqueous solution has the lowest vapor pressure.

A) 0.1 M KCl
B) 0.1 M Na₂CO₃
C) 0.2 M NaCl
D) 0.1 M MgCl₂
E) 0.2 M MgCl₂

Question 6

A solution contains 6.50 mol water, 0.300 mol sucrose, and 0.200 mol glucose. The solutes are nonvolatile. What is the vapor pressure of the solution at 35°C given that the vapor pressure of water is 42.2 torr?

A) 35.0 torr
B) 36.0 torr
C) 37.0 torr
D) 39.2 torr
E) 39.0 torr

Question 7

A solution is prepared by adding 1.50 mol glucose, which is not volatile, to 3.50 mol water. What is the vapor pressure of this solution at 25°C given that the vapor pressure of pure water is
23)8 torr?

A) 7.00 torr
B) 16.7 torr
C) 10.2 torr
D) 7.14 torr
E) 34.0 torr

Question 8

Which graph best describes how the vapor pressure of a solution varies according to Raoult's law as a nonvolatile solute is added to a liquid? The vapor pressure of the solution is plotted on the y-axis, and the mole fraction of solute is plotted on the x-axis. The origin (0, 0) is not necessarily located where the axes cross.

A) I
B) II
C) III
D) IV

Question 9

Indicate which aqueous solution has the slowest evaporation rate.

A) 0.1 M KCl
B) 0.2 M Na₂CO₃
C) 0.2 M NaCl
D) 0.1 M MgCl₂
E) 0.1 M NaBr

Question 10

Which intermolecular interactions are likely to result in the lowest vapor pressure of a substance?

A) dispersion interactions
B) dipole-dipole interactions
C) hydrogen bonds
D) ion-dipole interactions
E) ion-ion interactions

Question 11

Indicate which aqueous solution has the highest vapor pressure.

A) 0.1 M KCl
B) 0.2 M Na₂CO₃
C) 0.2 M NaCl
D) 0.1 M MgCl₂
E) 0.2 M MgCl₂

Question 12

Which statement below regarding vapor pressure is not correct?

A) Vapor pressure is an intensive property.
B) The substance with the stronger intermolecular forces has the lower vapor pressure.
C) Vapor pressure increases with increasing temperature.
D) Pure water has a higher vapor pressure at a given temperature than seawater.
E) A nonvolatile solute increases the vapor pressure of the solvent.

Question 13

A solution is prepared by adding 0.300 mol glucose, which is not volatile, to 4.50 mol water. What is the vapor pressure of this solution at 25°C given that the vapor pressure of pure water is
23)8 torr?

A) 9.38 torr
B) 22.3 torr
C) 23.4 torr
D) 1.49 torr
E) 22.5 torr

Question 14

In the diagram below, one beaker contains pure water and the other contains an equal volume of seawater. Seawater has various salts dissolved in it. The beakers are sitting in a totally enclosed chamber, and the outside temperature and pressure are held constant. Identify the statements below about this situation that are not correct.
I. Water will be transferred from the pure water beaker to the seawater beaker.
II. Water will be transferred from the seawater beaker to the pure water beaker.
III. The vapor pressure of the pure water is higher than the vapor pressure of the seawater.
IV. Pure water evaporates at a faster rate than seawater.
V. Water in the gas phase condenses into both beakers at the same rate.

A) II only
B) II and V
C) II, IV, and V
D) II, III, and IV
E) I, III, and IV

Question 15

Which graph best describes how the vapor pressure of a solution varies according to Raoult's law as a nonvolatile solute is added to a liquid? The vapor pressure of the solution is plotted on the y-axis, and the mole fraction of solvent is plotted on the x-axis. The origin (0, 0) is not necessarily located where the axes cross.

A) I
B) II
C) III
D) IV

Question 16

Which statement about how the vapor pressure (P) of a liquid depends on temperature is correct?

A) The larger the enthalpy of vaporization, the smaller the change in vapor pressure for a given change in temperature.
B) A graph of P versus 1/T is linear with a slope - $\Delta$ H_vap /R.
C) The smaller the intermolecular interactions, the greater the change in vapor pressure for a given change in temperature.
D) A graph of ln(P) versus T is linear with a slope - $\Delta$ H_vap /R.
E) A graph of ln(P) versus 1/T is linear with a slope - $\Delta$ H_vap /R.

Question 17

Which of the solutions shown here will have the lowest vapor pressure? White circles indicate solvent molecules; black circles indicate molecules of a nonvolatile solute.

A)
B)
C)
D)

Question 18

Which statement about the vapor pressure of a liquid is not correct?

A) The vapor pressure increases with increasing temperature because more molecules have sufficient energy to overcome the attractive forces that hold them in the liquid.
B) The vapor pressure increases with increasing strength of intermolecular interactions because molecules are better able to pull each other into the gas phase.
C) The pressure at which the rate of molecules evaporating equals the rate of molecules condensing is called the vapor pressure.
D) Adding a nonvolatile solute to a liquid decreases the vapor pressure.
E) The enthalpy of vaporization determines how sensitive the vapor pressure of a substance is to a change in temperature.

Question 19

In the diagram below, one beaker contains pure water and the other contains an equal volume of seawater. Seawater has various salts dissolved in it. The beakers are sitting in a totally enclosed chamber, and the outside temperature and pressure are held constant. Identify the statements below about this situation that are not correct. I. Water will be transferred from the pure water beaker to the seawater beaker.
II) The vapor pressure of the seawater is higher than the vapor pressure of the pure water.
III) Pure water evaporates at a faster rate than seawater.
IV) Water in the gas phase condenses into the seawater beaker at a faster rate than into the pure water beaker.

A) I, III, and IV
B) II, III, and IV
C) II and III
D) II and IV
E) II only

Question 20

Which of the solutions shown here will have the highest vapor pressure? White circles indicate solvent molecules; black circles indicate molecules of a nonvolatile solute.

A)
B)
C)
D)

Question 21

The vapor pressure of an aqueous solution is found to be 24.9 mm Hg at 25°C. What is the mole fraction of solute in this solution? The vapor pressure of water is 25.756 mm Hg at 25°C.

A) 0.967
B) 0.0332
C) 1.03
D) 0.0344
E) 0.976

Question 22

Which of the following gases do you expect to be most soluble in water?

A) CO₂
B) Ar
C) O₂
D) N₂
E) NH₃

Question 23

At 25°C, the vapor pressure of pure water is 25.756 mm Hg. What is the vapor pressure of water in a 0.500 m solution of sodium chloride?

A) 25.52 mm Hg
B) 0.2301 mm Hg
C) 0.4602 mm Hg
D) 25.35 mm Hg
E) 12.88 mm Hg

Question 24

Henry's law constant (mol/L · atm) for oxygen dissolving in blood is 3.74 * 10^-2 mol/(L atm) at body temperature, 37°C. Calculate the molar concentration of oxygen in blood for a scuba diver where the air pressure is 2.0 atm. The mole fraction of oxygen in air is 0.209.

A) 7.8 *10^-3 M
B) 2.7 *10^-3 M
C) 1.6 *10^-2 M
D) 1.3 * 10^-2 M
E) 0.11 M

Question 25

Which statement A-D about ionic solids is not correct?

A) The enthalpy of hydration includes two terms: one for ion-dipole interactions and one for dipole-dipole interactions.
B) The melting point of an ionic solid increases as the lattice energy increases.
C) The enthalpy of solution consists of three terms: one for the lattice energy, one for ion-dipole interactions, and one for dipole-dipole interactions.
D) The lattice energy depends inversely on the product of ionic charges, directly on the distance between the ions, and on the structure of the solid.
E) Statements A-D are all correct.

Question 26

Which of the following must be true if the dissolution of an ionic solid in water is endothermic?

A) The change in enthalpy for this process is negative.
B) The process releases energy.
C) The hydration energy exceeds the lattice energy.
D) The ions have charges greater than 1 unit.
E) The solubility will increase with increasing temperature.

Question 27

Air (consisting mostly of nitrogen (N₂), oxygen (O₂), and argon (Ar)) dissolves in water due to ________

A) dispersion forces.
B) dipole-induced dipole forces.
C) ion-dipole forces.
D) hydrogen bonding.
E) dipole-dipole forces.

Question 28

The pressure of carbon dioxide in a beer bottle is about 2 atm. At this pressure about 0.15 g of carbon dioxide dissolves in 100 mL of beer. How much carbon dioxide remains dissolved in 100 mL after the bottle is opened? The partial pressure of carbon dioxide is 2.0 *10^-4 atm.

A) 0.15 g
B) 3.8 *10^-3 g
C) 6.0 * 10^-5 g
D) 3.0 * 10^-4 g
E) 1.5 * 10^-5 g

Question 29

Which statement A-D about ionic solids is not correct?

A) The enthalpy of hydration includes two terms: one for ion-dipole interactions and one for dipole-dipole interactions.
B) The melting point of an ionic solid increases as the lattice energy increases.
C) The enthalpy of solution consists of two terms: the lattice energy and the enthalpy of hydration.
D) The lattice energy depends directly on the product of ionic charges, inversely on the distance between the ions, and on the structure of the solid.
E) Statements A-D are all correct.

Question 30

Molecular nitrogen (N₂) interacts with water and is sparingly soluble in water due to ________

A) dispersion forces.
B) dipole-induced dipole forces.
C) ion-dipole forces.
D) hydrogen bonding.
E) dipole-dipole forces.

Question 31

Indicate which aqueous solution has the fastest evaporation rate.

A) 0.1 M KCl
B) 0.2 M Na₂CO₃
C) 0.2 M NaCl
D) 0.1 M MgCl₂
E) 0.2 M MgCl₂

Question 32

The solubility of any gas in a liquid can always be increased by ________

A) increasing the temperature.
B) using a solvent that can form hydrogen bonds.
C) using a solvent that is nonpolar.
D) decreasing the pressure of the gas above the solvent.
E) increasing the pressure of the gas above the solvent.

Question 33

The pressure inside a bottle of carbonated beverage decreases when it is cooled in a refrigerator. What is the main reason for this change?

A) The solubility of carbon dioxide is higher at the lower temperature.
B) Some carbon dioxide condenses to form a liquid phase.
C) The volume of the vapor contracts.
D) The volume of the liquid contracts.
E) The volume of the bottle contracts.

Question 34

The carbon dioxide pressure in a bottle of champagne is about 6 atm. At this pressure about 0.45 g of carbon dioxide dissolves in 100 mL of champagne. How much carbon dioxide remains dissolved in 100 mL after the bottle is opened? The partial pressure of carbon dioxide then is 6.0 *10^-4 atm.

A) 0.45 g
B) 3.8 *10^-3 g
C) 1.7 *10^-6 g
D) 9.0 * 10^-4 g
E) 4.5 *10^-5 g

Question 35

Henry's law constant (mol/L · atm) for oxygen dissolving in blood is 3.74 *10^-2 mol/L · atm at body temperature, 37°C. Calculate the molar concentration of oxygen in blood for an alpine climber where the atmospheric pressure is 0.45 atm. The mole fraction of oxygen in air is 0.209.

A) 7.8 *10^-3 M
B) 3.5 *10^-3 M
C) 2.3 *10^-2 M
D) 1.3 *10^-2 M
E) 0.11 M

Question 36

Predict the relationship between the lattice energies (U) of NaCl and MgO.

A) U_MgO > 4 U_NaCl
B) U_MgO < 4 U_NaCl
C) U_MgO > 2 U_NaCl
D) U_MgO < 2 U_NaCl
E) U_MgO and U_NaCl are about the same.

Question 37

Which arrangement below orders the cations from smallest to largest hydration energy? Smallest to largest means least negative hydration energy to most negative hydration energy.

A) Li⁺ < Na⁺ < Ca²⁺ < Mg²⁺
B) Na⁺ < Mg²⁺ < Ca²⁺ < Li⁺
C) Na⁺ < Li⁺ < Ca²⁺ < Mg²⁺
D) Na⁺ < Mg²⁺ < Li⁺ < Ca²⁺
E) Li⁺ < Na⁺ < Mg²⁺ < Ca²⁺

Question 38

Predict the relationship between the lattice energies (U) of KBr and CaO.

A) U_CaO > 4 U_KBr because the ionic charges are +2 and -2 and the distance between the ions is smaller.
B) U_CaO < 4 U_KBr because the ionic charges are +2 and -2 and the distance between the ions is larger.
C) U_CaO > 2 U_KBr because the ionic charges are +2 and -2 and the distance between the ions is smaller.
D) U_CaO < 4 U_KBr because the ionic charges are +2 and -2 and the distance between the ions is larger..
E) U_CaO and U_KBr are about the same because the larger ionic charge is compensated by the larger distance between the ions.

Question 39

Coulomb's law states that the interaction energy between ions depends ________

A) only on the ionic charges.
B) only on the distance between the ions.
C) directly on both the ionic charges and the distance between the ions.
D) on the temperature.
E) directly on the ionic charges and inversely on the distance between the ions.

Question 40

If water contains about 42 mg of oxygen per liter at 20°C and 1.0 atm, what would be the value of Henry's law constant for oxygen dissolving in water? The mole fraction of oxygen in air is 0.21.

A) 0.25 mol/(L atm)
B) 1.3 *10^-3 mol/(L atm)
C) 1.9 *10^-2 mol/(L atm)
D) 6.3 * 10^-3 mol/(L atm)
E) 0.010 mol/(L atm)

Question 41

Which of the following ranks the compounds from lowest to highest lattice energy?

A) NaF < MgF₂ < CaF₂ < KF
B) KF < NaF < CaF₂ < MgF₂
C) MgF₂ < CaF₂ < NaF < KF
D) CaF₂ < KF < NaF < MgF₂
E) MgF₂ < CaF₂ < KF < NaF

Question 42

Which of the following is needed to calculate the lattice energy of an ionic compound?

A) enthalpy of solution
B) enthalpy of combustion
C) specific heat
D) ionization energy
E) enthalpy of solvation

Question 43

Which listing of ionic compounds is in order of increasing melting point?

A) NaF < MgF₂ < AlF₃
B) MgF₂ < NaF < AlF₃
C) AlF₃ < MgF₂ < NaF
D) AlF₃ < NaF < MgF₂
E) NaF < AlF₃ < MgF₂

Question 44

Which of the following requires the smallest energy to separate the ions?

A) CaF₂
B) KF
C) NaF
D) MgF₂
E) LiF

Question 45

Which one of the ionic compounds below would you expect to have the largest (most negative) lattice energy?

A) MgF₂
B) MgCl₂
C) MgBr₂
D) MgI₂
E) CaI₂

Question 46

Use the following data to calculate the enthalpy change for the following reaction: $$\begin{array} { l c } { \mathrm { K } ( s ) + \frac { 1 } { 2 } \mathrm { Cl } _ { 2 } ( g ) \rightarrow \mathrm { KCl } ( s ) } & \\ { \text { Quantity } } & \text { Magnitude } ( \mathbf { k J } / \mathbf { m o l } ) \\\text { Ionization energy of } \mathrm { K } ( g ) & 425 \\\text { Electron affinity of } \mathrm { Cl } ( g ) & - 349 \\\text { Vaporization energy of } \mathrm { K } ( s ) & 89 \\\text { Bond energy of } \mathrm { Cl } _ { 2 } ( g ) & 240 \\\text { Lattice energy of } \mathrm { KCl } ( s ) & - 723\end{array}$$

A) -603 kJ/mol
B) -11.0 kJ/mol
C) 405 kJ/mol
D) -318 kJ/mol
E) -438 kJ/mol

Question 47

Use the following data to calculate the enthalpy change for the following reaction. $$\begin{array} { l c } { 2 \mathrm { Na } ( s ) + \frac { 1 } { 2 } \mathrm { O } _ { 2 } ( g ) \rightarrow \mathrm { Na } _ { 2 } \mathrm { O } ( s ) } & \\ { \text { Quantity } } & \text { Magnitude } ( \mathbf { k J } / \mathbf { m o l } ) \\\text { Ionization energy of } \mathrm { Na } ( \mathrm { g } ) & 495 \\\text { Electron affinity of } \mathrm { O } ( g ) \text { for } 2 \mathrm { e } & 603 \\\text { Vaporization energy of } \mathrm { Na } ( s ) & 109 \\\text { Bond energy of } \mathrm { O } _ { 2 } ( g ) & 499 \\\text { Lattice energy for } \mathrm { Na } _ { 2 } \mathrm { O } ( s ) & - 2,477\end{array}$$

A) -1,322 kJ/mol
B) -1,571 kJ/mol
C) -557 kJ/mol
D) -771 kJ/mol
E) -416 kJ/mol

Question 48

Why would you expect or predict sodium fluoride to have the lowest solubility in water compared to the other sodium halide salts?

A) Fluoride has the largest (most negative) enthalpy of hydration of all the halide ions.
B) Fluoride has the smallest (least negative) enthalpy of hydration of all the halide ions.
C) Sodium fluoride has the largest (most negative) lattice energy of all the sodium halides.
D) Sodium fluoride has the smallest molar mass of all the sodium halides.
E) Fluoride is the most basic of all the halide ions and forms HF in aqueous solution.

Question 49

Which of the following will have the largest lattice energy?

A) NaF
B) NaCl
C) NaBr
D) NaI
E) CsCl

Question 50

Which of the following will require the greatest energy input to separate the ions?

A) MgI₂
B) MgF₂
C) MgCl₂
D) MgBr₂
E) NaCl

Question 51

Use the following data to calculate the lattice energy of lithium oxide, Li₂O. $$\begin{array} { l c } { \text { Quantity } } & \text { Magnitude } ( \mathrm { kJ } / \mathbf { m o l } ) \\\text { Ionization energy of } \mathrm { Li } ( g ) & 519 \\\text { Electron affinity of } \mathrm { O } ( g ) \text { for } 2 \mathrm { e } & 603 \\\text { Vaporization energy of } \mathrm { Li } ( s ) & 147 \\\text { Bond energy of } \mathrm { O } _ { 2 } ( g ) & 499 \\\text { Reaction enthalpy: } 2 \mathrm { Li } ( s ) + \frac { 1 } { 2 } \mathrm { O } _ { 2 } ( g ) \rightarrow \mathrm { Li } _ { 2 } \mathrm { O } ( s ) & - 20\end{array}$$

A) -2,586 kJ/mol
B) -2,972 kJ/mol
C) -3,081 kJ/mol
D) -2,205 kJ/mol
E) -2,831 kJ/mol

Question 52

Use the following data to calculate the lattice energy of lithium chloride, LiCl. $$\begin{array} { l c } { \text { Quantity } } & \text { Magnitude } ( \mathbf { k J } / \mathbf { m o l } ) \\\text { Ionization energy of } \mathrm { Li } ( g ) & 519 \\\text { Electron affinity of } \mathrm { Cl } ( g ) & - 349 \\\text { Vaporization energy of } \mathrm { Li } ( s ) & 147 \\\text { Bond energy of } \mathrm { Cl } _ { 2 } ( g ) & 240 \\\text { Reaction enthalpy: } \mathrm{Li}(s)+\frac{1}{2} \mathrm{Cl}_{2}(g) \rightarrow \mathrm{LiCl}(s)& - 409\end{array}$$

A) -603 kJ/mol
B) -11.0 kJ/mol
C) +405 kJ/mol
D) -846 kJ/mol
E) -131 kJ/mol

Question 53

Which of the following compounds has the largest melting point?

A) NaF
B) KCl
C) RbCl
D) BeF₂
E) NaCl

Question 54

What physical property is used to separate the hydrocarbon components in petroleum (crude oil)?

A) melting point
B) density
C) boiling point
D) molar mass
E) viscosity

Question 55

Calculate the lattice energy of sodium fluoride from the following data:
Ionization energy of Na: 496 kJ/mol
Electron affinity of F: -328 kJ/mol
Energy to vaporize Na: 108 kJ/mol
F₂ bond energy: 160 kJ/mol
Energy change for the reaction:
Na(s) + $\frac { 1 } { 2 }$
F₂(g) $\rightarrow$ NaF(s); $\Delta$ H = -575 kJ

A) 931 kJ/mol
B) -931 kJ/mol
C) -1011 kJ/mol
D) 1,011 kJ/mol
E) -851 kJ/mol

Question 56

Determine the energy change for the reaction
Li(s)+ $\frac { 1 } { 2 }$ Cl₂(g) $\rightarrow$ LiCl(s)
From the following data:
Lattice energy of LiCl = -861 kJ/mol
Energy to vaporize Li = 159 kJ/mol
Ionization energy of Li = 520 kJ/mol
Cl₂ bond energy: 240 kJ/mol
Electron affinity of Cl: -349 kJ/mol

A) -411 kJ/mol
B) -528 kJ/mol
C) 311 kJ/mol
D) -861 kJ/mol
E) -291 kJ/mol

Question 57

Which one of the ionic compounds below would you expect to have the largest (most negative) lattice energy?

A) MgO
B) CaO
C) SrO
D) BaO
E) BeO

Question 58

Which of the following is not typically needed to calculate the lattice energy of an ionic compound?

A) enthalpy of vaporization
B) bond enthalpy
C) enthalpy of solution
D) electron affinity
E) ionization energy

Question 59

Which one of the ionic compounds below would you expect to have the smallest (least negative) lattice energy?

A) MgF₂
B) MgCl₂
C) MgBr₂
D) MgI₂
E) CaI₂

Question 60

Which one of the ionic compounds below would you expect to have the smallest (least negative) lattice energy?

A) MgO
B) CaO
C) SrO
D) BaO
E) BeO

Question 61

A solution is prepared by mixing 75 g of methanol (CH₃OH) with 25 g of ethanol (CH₃CH₂OH). Use the following data to determine the vapor pressure of this solution at 20°C. $\begin{array}{cc} \text { Substance }&\text { Vapor Pressure }\\& \text { at } \mathbf{2 0}{ }^{\circ} \mathbf{C} \text { (torr) } \\\text { Methanol } & 92 \\\text { Ethanol } & 45\end{array}$

A) 69 torr
B) 57 torr
C) 80 torr
D) 73 torr
E) 83 torr

Question 62

What is the molality of a solution produced by dissolving 14.40 g of LiCl (42.39 g/mol) in water to make 0.104 L of solution with a density of 1.102 g/mL?

A) 0.340 m
B) 3.39 m
C) 3.27 m
D) 3.74 m
E) 2.96 m

Question 63

What would be the freezing point of a 1 m solution of HCl in water (K_f = 1.86°C)?

A) 0.0°C
B) -1.86°C
C) -3.72°C
D) 1.86°C
E) -0.93°C

Question 64

Calculate the molality of a solution containing 0.755 mol glucose and 1.75 kg of water.

A) 0.875 m
B) 0.583 m
C) 0.431 m
D) 580 m
E) 0.850 m

Question 65

Which of the following aqueous solutions will have the lowest freezing point?

A) 0.1 m magnesium sulfate, MgSO₄
B) 0.1 m potassium chloride, KCl
C) 0.08 m magnesium chloride, MgCl₂
D) 0.04 m sodium sulfate, Na₂SO₄
E) 0.05 m sodium chloride, NaCl

Question 66

How many moles of solute are there in a 0.155 m glucose solution prepared with 50.0 kg of water?

A) 15.5 mol
B) 15.0 mol
C) 0.155 mol
D) 31.0 mol
E) 7.75 mol

Question 67

A solution is prepared by mixing 50 g of methanol (CH₃OH) with 50 g of ethanol (CH₃CH₂OH). Use the following data to determine the vapor pressure of this solution at 20°C. $\begin{array}{cc} \text { Substance }&\text { Vapor Pressure }\\& \text { at } \mathbf{2 0 ^ { \circ }} \mathbf{C} \text { (torr) } \\\text { Methanol } & 92 \\\text { Ethanol } & 45\end{array}$

A) 69 torr
B) 57 torr
C) 80 torr
D) 73 torr
E) 83 torr

Question 68

Determine the molal concentration of a sugar solution in water that has a freezing point of -2.1°C. K_f = 1.86°C/m for water.

A) 1.13 m
B) -1.13 m
C) 3.91 m
D) -3.91 m
E) 0.113 m

Question 69

Isopropyl alcohol has a boiling point of 82.3°C. Solutions of isopropyl alcohol in water have normal boiling points less than 100°C. Is this observation consistent with the following equation? Explain.
$\Delta$ T_b = iK_bm

A) No, boiling point elevation only pertains to ionic solutes.
B) No, boiling point elevation only pertains to nonvolatile solutes.
C) Yes, boiling point depression occurs for all solutions.
D) Yes, the boiling point depression constant for isopropyl alcohol is large.
E) No, isopropyl alcohol can hydrogen bond with water.

Question 70

A petroleum company separates benzene (78 g/mol) and n-octane (114 g/mol) by distillation. In one run, the starting solution was 15% benzene and 85% n-octane by mass at 25°C. In the liquid, the mole ratio of benzene to n-octane was ________, and in the vapor it was ________. The vapor pressures for benzene and n-octane at this temperature are 11 torr and 95 torr, respectively.

A) 0.26; 0.12
B) 0.18; 0.12
C) 0.80; 0.20
D) 0.26; 0.18
E) 0.31; 0.23

Question 71

The concentration unit of molality is symbolized as

A) M
B) m
C) M..
D) mol
E) mo

Question 72

A solution is made by dissolving 100 g of essentially nonvolatile ethylene glycol (C₂H₆O₂) in
500 g of water. What is the resulting freezing point of the solution (K_f = 1.86°C/m)?

A) -5.99°C
B) -1.86°C
C) -0.372°C
D) -1.61°C
E) -3.23°C

Question 73

A medical company separates methanol (CH₃OH)) and ethanol (CH₃CH₂OH) by distillation. In one run, the starting solution was 25% methanol and 75% ethanol by mass at 25°C. In the liquid, the mole ratio of methanol to ethanol was ________, and in the vapor it was ________. The vapor pressures for methanol and ethanol at this temperature are 92 torr and 45 torr, respectively.

A) 0.48; 2.0
B) 2.0; 0.48
C) 0.33; 2.0
D) 0.48; 0.33
E) 2.3; 0.55

Question 74

Calculate the molality of a solution containing 0.355 mol sodium hydrogen carbonate (baking soda) and 245 g of water.

A) 1.45 m
B) 1.12 *10^-2 m
C) 2.11 m
D) 1.12 *10^-3 m
E) 0.0211 m

Question 75

The normal temperature range of the liquid phase of pure water is 0°C to 100°C. Which of the following solutions will have the largest temperature range for the liquid state?

A) 1 M aqueous ethanol solution
B) 1 M aqueous potassium bromide solution
C) 1 M aqueous acetic acid solution
D) 1 M aqueous magnesium bromide solution
E) 1 M aqueous magnesium sulfate

Question 76

What mass of a 0.66 m ammonium nitrate (80.04 g/mol) solution will contain 0.15 mol of solute?

A) 227 g
B) 239 g
C) 440 g
D) 493 g
E) 4,453 g

Question 77

Which of the following is a colligative property?
I. osmotic pressure
II. vapor pressure
III. density
IV. freezing point depression
V. boiling point

A) I, II, IV, and V only
B) I and IV only
C) II and IV only
D) II, III, and V only
E) I, II, and V only

Question 78

The molality of a solution is defined as ________

A) moles of solute per liter of solution.
B) grams of solute per liter of solution.
C) moles of solute per kilogram of solution.
D) moles of solute per kilogram of solvent.
E) grams of solute per kilogram of solvent.

Question 79

Which of the following is a colligative property?
I. osmotic pressure
II. vapor pressure lowering
III. freezing point depression
IV. boiling point elevation

A) II, III, and IV only
B) III and IV only
C) II and III only
D) I, III, and IV only
E) I, II, III, and IV

Question 80

Which statement regarding the fractional distillation of two substances is not correct?

A) The substance with the lower boiling point has a higher concentration in the vapor than in the liquid.
B) The boiling point of the mixture is not constant, changing as the distillation progresses.
C) The substance with the higher vapor pressure has a higher concentration in the vapor than in the liquid.
D) The mole ratio of the two substances in the vapor is not the same as it is in the liquid.
E) When the vapor at any stage of the distillation is condensed, the resulting distillate has the same composition as the original mixture of the two substances.