Chapter 10

The graphs in Figure \(\mathrm{P} 10.5\) have the same scales and describe the change in \(\ln \left(P_{\text {vap }}\right)\) of two pure liquids as a function of temperature. Which liquid has the stronger intermolecular attractive forces? Explain your answer.

Which of the drawings in Figure P10.6, both of which are at constant temperature, most likely illustrates the pure liquid with the lower normal boiling point? Explain your choice.

Which type of intermolecular force exists in all substances?

At room temperature, bromine \(\left(\mathrm{Br}_{2}\right)\) is a corrosive red liquid, whereas iodine \(\left(\mathrm{I}_{2}\right)\) is a volatile violet solid. The differences point to different strengths of intermolecular forces between these halogens, with those for \(\mathrm{I}_{2}\) being stronger. What kind of intermolecular force is responsible for these differences?

Why do gases behave nonideally at high pressures and low temperatures?

Why are normal boiling points generally lower for branched hydrocarbons than for straight-chain hydrocarbons of the same molecular mass?

In each of the following pairs of molecules, which one experiences the stronger dispersion forces? (a) \(\mathrm{CCl}_{4}\) or \(\mathrm{CF}_{4};\) (b) \(\mathrm{CH}_{4}\) or \(\mathrm{C}_{3} \mathrm{H}_{8}\)

What kinds of intermolecular forces must be overcome as solid CO\(_{2}\) sublimes?

How are the water molecules preferentially oriented around the anion in an aqueous solution of sodium chloride?

How are the water molecules preferentially oriented around the cation in an aqueous solution of potassium bromide?

Why are dipole-dipole interactions generally weaker than ion-dipole interactions?

Two liquids-one polar, one nonpolar-have the same molar mass. Which one is likely to have the higher boiling point? Explain your answer.

Why are hydrogen bonds considered a special class of dipole-dipole interactions?

Can all polar hydrogen-containing molecules form hydrogen bonds? Why or why not?

The permanent dipole moment of \(\mathrm{CH}_{2} \mathrm{F}_{2}(1.93 \mathrm{D})\) is larger than that of \(\mathrm{CH}_{2} \mathrm{Cl}_{2}(1.60 \mathrm{D}),\) yet the boiling point of \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\left(40^{\circ} \mathrm{C}\right)\) is much higher than that of \(\mathrm{CH}_{2} \mathrm{F}_{2}\) \(\left(-52^{\circ} \mathrm{C}\right) .\) Why?

How is it that the permanent dipole moment of HCl \((1.08 \mathrm{D})\) is larger than the permanent dipole moment of HBr \((0.82 \mathrm{D}),\) yet HBr boils at a higher temperature?

Hydrogen peroxide \(\left(\mathrm{H}_{2} \mathrm{O}_{2}\right)\) and water \(\left(\mathrm{H}_{2} \mathrm{O}\right)\) are both liquids at room temperature, but their standard heats of vaporization \(\Delta H_{\mathrm{vap}}^{o}\) are different: \(52 \mathrm{kJ} / \mathrm{mol}\) and \(41 \mathrm{kJ} / \mathrm{mol},\) respectively. Which substance has the stronger intermolecular forces? Can you suggest why?

Explain why the melting point of methyl fluoride, \(\mathrm{CH}_{3} \mathrm{F}\) \(\left(-142^{\circ} \mathrm{C}\right),\) is higher than the melting point of methane, \(\mathrm{CH}_{4}\left(-182^{\circ} \mathrm{C}\right)\).

Explain why the boiling point of \(\mathrm{Br}_{2}\left(59^{\circ} \mathrm{C}\right)\) is lower than that of iodine monochloride, ICl \(\left(97^{\circ} \mathrm{C}\right)\), even though they have nearly the same molar mass.

Why doesn't fluoromethane (CH \(_{3} \mathrm{F}\) ) exhibit hydrogen bonding, whereas hydrogen fluoride, HF, does?

The boiling point of phosphine, \(\mathrm{PH}_{3}\left(-88^{\circ} \mathrm{C}\right),\) is lower than that of ammonia, \(\mathrm{NH}_{3}\left(-33^{\circ} \mathrm{C}\right)\), even though \(\mathrm{PH}_{3}\) has twice the molar mass of \(\mathrm{NH}_{3} .\) Why?

In which of the following compounds do the molecules experience the strongest dipole-dipole attractions? (a) \(\mathrm{CF}_{4}\) (b) \(\mathrm{CF}_{2} \mathrm{Cl}_{2} ;\) (c) \(\mathrm{CCl}_{4}\)

Which of the following compounds, \(\mathrm{CO}_{2}, \mathrm{NO}_{2}, \mathrm{SO}_{2},\) or \(\mathrm{H}_{2} \mathrm{S},\) is expected to have the weakest interactions between its molecules?

Which of the following molecules can form hydrogen bonds among themselves in pure samples of bulk material? (a) methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right) ;\) (b) ethane \(\left(\mathrm{CH}_{3} \mathrm{CH}_{3}\right) ;(\mathrm{c})\) dimethyl ether \(\left(\mathrm{CH}_{3} \mathrm{OCH}_{3}\right) ;\) (d) acctic acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\)

Which of the following molecules can form hydrogen bonds with molecules of water? (a) methanol (CH \(_{3} \mathrm{OH}\) ); (b) ethane \(\left(\mathrm{CH}_{3} \mathrm{CH}_{3}\right) ;\) (c) dimethyl ether \(\left(\mathrm{CH}_{3} \mathrm{OCH}_{3}\right)\) (d) acetic acid (CH \(_{3} \mathrm{COOH}\) )

What is the difference between the terms miscible and insoluble?

What properties of water molecules enable them to hydrate and separate cations and anions in aqueous solutions?

One of the compounds in Figure \(\mathrm{P} 10.16\) is insoluble in water; the other has a water solubility of \(0.87 \mathrm{g} / 100 \mathrm{mL}\) at \(20^{\circ} \mathrm{C} .\) Identify which is which, and explain your reasoning.

Explain why the solubility of the series of alcohols with the formula \(\mathrm{C}_{n} \mathrm{H}_{2 n+2} \mathrm{OH}\) decreases with increasing \(n\).

How does the presence of increasingly longer hydrocarbon chains in the structure affect the solubility of a series of structurally related molecules in water?

In each of the following pairs of compounds, which compound is likely to be more soluble in water? a. \(\mathrm{CCl}_{4}\) or \(\mathrm{CHCl}_{3}\) b. \(\mathrm{CH}_{3} \mathrm{OH}\) or \(\mathrm{C}_{6} \mathrm{H}_{11} \mathrm{OH}\) c. NaF or MgO d. \(\mathrm{CaF}_{2}\) or \(\mathrm{BaF}_{2}\)

In each of the following pairs of compounds, which compound is likely to be more soluble in \(\mathrm{CCl}_{4} ?\) a. \(\mathrm{Br}_{2}\) or \(\mathrm{NaBr}\) b. \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\) or \(\mathrm{CH}_{3} \mathrm{OCH}_{3}\) c. \(\mathrm{CS}_{2}\) or \(\mathrm{KOH}\) d. \(I_{2}\) or \(C a F_{2}\)

Which of these pairs of substances is likely to be miscible? a. \(\mathrm{Br}_{2}\) and benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)\) b. \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OCH}_{2} \mathrm{CH}_{3}\) (diethyl ether) and \(\mathrm{CH}_{3} \mathrm{COOH}\) (acetic acid) c. \(\mathrm{C}_{6} \mathrm{H}_{12}\) (cyclohexane) and hexane \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}\right)\) d. \(\mathrm{CS}_{2}\) (carbon disulfide) and \(\mathrm{CCl}_{4}\) (carbon tetrachloride)

Which of these pairs of substances is likely to be miscible? a. \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\) (ethanol) and \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OCH}_{2} \mathrm{CH}_{3}\) (diethyl ether) b. \(\mathrm{CH}_{3} \mathrm{OH}\) (methanol) and methyl amine \(\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)\) c. CH \(_{3} \mathrm{CN}\) (acetonitrile) and acetone (CH \(_{3} \mathrm{COCH}_{3}\) ) d. \(\mathrm{CF}_{3} \mathrm{CHF}_{2}\) (a Freon replacement) and \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}\) (pentane)

Which of the following compounds is likely to be the most soluble in water? (a) \(\mathrm{NaCl} ;\) (b) \(\mathrm{KI} ;\) (c) \(\mathrm{Ca}(\mathrm{OH})_{2} ;\) (d) \(\mathrm{CaO}\).

Which of these substances is the least soluble in water? a. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{2} \mathrm{CH}_{2} \mathrm{OH}\) b. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{4} \mathrm{CH}_{2} \mathrm{OH}\) c. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{6} \mathrm{CH}_{2} \mathrm{OH}\) d. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{8} \mathrm{CH}_{2} \mathrm{OH}\)

Which of these substances is the most soluble in water? a. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}\) b. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{4} \mathrm{CH}_{2} \mathrm{Cl}\) c. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{6} \mathrm{CH}_{2} \mathrm{Br}\) d. \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{8} \mathrm{CH}_{2} \mathrm{I}\)

Why does the solubility of most gases in most liquids increase with decreasing temperature?

Which term, \(k_{\mathrm{H}}\) or \(P\), in Henry's law is affected by temperature?

Air is primarily a mixture of nitrogen and oxygen. Is the Henry's law constant for the solubility of air in water the sum of \(k_{\mathrm{H}}\) for \(\mathrm{N}_{2}\) and \(k_{\mathrm{H}}\) for \(\mathrm{O}_{2} ?\) Explain why or why not.

Why is the Henry's law constant for \(\mathrm{CO}_{2}\) so much larger than those for \(\mathrm{N}_{2}\) and \(\mathrm{O}_{2}\) at the same temperature?

Which sulfur oxide would you predict to be more soluble in nonpolar solvents, \(\mathrm{SO}_{2}\) or \(\mathrm{SO}_{3} ?\)

Which oxide would you predict to be more soluble in polar solvents, \(\mathrm{CO}_{2}\) or \(\mathrm{NO}_{2} ?\)

The solubility of \(\mathrm{O}_{2}\) in water is \(6.5 \mathrm{mg} / \mathrm{L}\) at an atmospheric pressure of 1 atm and temperature of \(40^{\circ} \mathrm{C}\). Calculate the Henry's law constant of \(\mathrm{O}_{2}\) at \(40^{\circ} \mathrm{C}\). The mole fraction of \(\mathrm{O}_{2}\) in air is 0.209.

The solubility of air in water is approximately \(7.9 \times 10^{-4} \mathrm{M}\) at \(20^{\circ} \mathrm{C}\) and 1.0 atm. Calculate the Henry's law constant for air. Is the \(k_{\mathrm{H}}\) value of air approximately equal to the sum of the \(k_{\mathrm{H}}\) values for \(\mathrm{N}_{2}\) and \(\mathrm{O}_{2}\) because these two gases make up \(99 \%\) of the gases in air?

Manufacturers of carbonated beverages dissolve \(\mathrm{CO}_{2}(g)\) in water under pressure to produce sodas. If a manufacturer uses a pressure of 4.16 atm at \(25^{\circ} \mathrm{C}\) to carbonate the water, how many grams of \(\mathrm{CO}_{2}\) are in the average can of soda, the volume of which is 355 mL? The Henry's law constant for \(\mathrm{CO}_{2}\) at \(25^{\circ} \mathrm{C}\) is \(0.03360 \mathrm{mol} /(\mathrm{L} \cdot \mathrm{atm})\).

Why does the vapor pressure of a liquid increase as temperature increases?

Which of the following factors influences the vapor pressure of a pure liquid? a. the volume of liquid present in a container b. the temperature of the liquid c. the surface area of the liquid

Is vapor pressure an intensive or extensive property of a liquid?

A chef observes bubbles while heating a pot of water to \(60^{\circ} \mathrm{C}\) to poach vegetables. What are the gases in the bubbles and where did they come from?