Chapter 11

List the three states of matter in order of (a) increasing molecular disorder and (b) increasing intermolecular attractions. (c) Which state of matter is most easily compressed?

(a) How does the average kinetic energy of molecules compare with the average energy of attraction between molecules in solids, liquids, and gases? (b) Why does increasing the temperature cause a solid substance to change in succession from a solid to a liquid to a gas? (c) What happens to a gas if you put it under extremely high pressure?

If you mix olive oil with water, the olive oil will float on top of the water. The density of water is \(1.00 \mathrm{~g} / \mathrm{cm}^{3}\) at room temperature. (a) Is the density of olive oil more or less than \(1.00 \mathrm{~g} / \mathrm{cm}^{3} ?\) (b) The density of olive oil in its liquid phase does vary with temperature. Do you think olive oil would be more dense or less dense at higher temperatures? Explain.

Benzoic acid, \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\), melts at \(122^{\circ} \mathrm{C}\). The density in the liquid state at \(130^{\circ} \mathrm{C}\) is \(1.08 \mathrm{~g} / \mathrm{cm}^{3} .\) The density of solid benzoic acid at \(15^{\circ} \mathrm{C}\) is \(1.266 \mathrm{~g} / \mathrm{cm}^{3} .\) (a) In which of these two states is the average distance between molecules greater? (b) Explain the difference in densities at the two temperatures in terms of the relative kinetic energies of the molecules.

Which type of intermolecular attractive force operates between (a) all molecules, (b) polar molecules, (c) the hydrogen atom of a polar bond and a nearby small electronegative atom?

Based on what you have learned about intermolecular forces, would you say that matter is fundamentally attracted or repulsed by other matter?

Describe the intermolecular forces that must be overcome to convert each of the following from a liquid or solid to a gas: (a) \(\mathrm{I}_{2}\), (b) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\), (c) \(\mathrm{H}_{2} \mathrm{Se}\).

What type of intermolecular force accounts for the following differences in each case? (a) \(\mathrm{CH}_{3} \mathrm{OH}\) boils at \(65^{\circ} \mathrm{C}, \mathrm{CH}_{3} \mathrm{SH}\) boils at \(6^{\circ} \mathrm{C}\). (b) Xe is liquid at atmos- pheric pressure and \(120 \mathrm{~K}\), whereas \(\mathrm{Ar}\) is a gas. (c) \(\mathrm{Kr}\), atomic weight 84 , boils at \(120.9 \mathrm{~K}\), whereas \(\mathrm{Cl}_{2}\), molecular weight about 71, boils at \(238 \mathrm{~K} .\) (d) Acetone boils at \(56{ }^{\circ} \mathrm{C}\), whereas 2 -methylpropane boils at \(-12{ }^{\circ} \mathrm{C}\).

(a) What is meant by the term polarizability? (b) Which of the following atoms would you expect to be most polarizable: \(\mathrm{N}, \mathrm{P}, \mathrm{As}, \mathrm{Sb}\) ? Explain. (c) Put the following molecules in order of increasing polarizability: \(\mathrm{GeCl}_{4}, \mathrm{CH}_{4}\) \(\mathrm{SiCl}_{4}, \mathrm{SiH}_{4}\), and \(\mathrm{GeBr}_{4}\). (d) Predict the order of boiling points of the substances in part (c).

True or false: (a) The more polarizable the molecules, the stronger the dispersion forces between them. (b) The boiling points of the noble gases decrease as you go down the column in the periodic table. (c) In general, the smaller the molecule, the stronger the dispersion forces. (d) All other factors being the same, dispersion forces between molecules increase with the number of electrons in the molecules.

Which member of the following pairs has the larger London dispersion forces: (a) \(\mathrm{H}_{2} \mathrm{O}\) or \(\mathrm{H}_{2} \mathrm{~S},(\mathrm{~b}) \mathrm{CO}_{2}\) or \(\mathrm{CO},(\mathrm{c}) \mathrm{SiH}_{4}\) or \(\mathrm{GeH}_{4} ?\)

Butane and 2 -methylpropane, whose space-filling models are shown, are both nonpolar and have the same molecular formula, yet butane has the higher boiling point \(\left(-0.5^{\circ} \mathrm{C}\right.\) compared to \(\left.-11.7{ }^{\circ} \mathrm{C}\right) .\) Explain.

(a) What atoms must a molecule contain to participate in hydrogen bonding with other molecules of the same kind? (b) Which of the following molecules can form hydrogen bonds with other molecules of the same kind: \(\mathrm{CH}_{3} \mathrm{~F}, \mathrm{CH}_{3} \mathrm{NH}_{2}, \mathrm{CH}_{3} \mathrm{OH}, \mathrm{CH}_{3} \mathrm{Br} ?\)

Rationalize the difference in boiling points between the members of the following pairs of substances: (a) HF \(\left(20^{\circ} \mathrm{C}\right)\) and \(\mathrm{HCl}\left(-85{ }^{\circ} \mathrm{C}\right)\), (b) \(\mathrm{CHCl}_{3}\left(61{ }^{\circ} \mathrm{C}\right)\) and \(\mathrm{CHBr}_{3}\) \(\left(150^{\circ} \mathrm{C}\right)\) (c) \(\mathrm{Br}_{2}\left(59^{\circ} \mathrm{C}\right)\) and \(\mathrm{ICl}\left(97{ }^{\circ} \mathrm{C}\right)\)

Identify the types of intermolecular forces present in each of the following substances, and select the substance in each pair that has the higher boiling point: (b) \(\mathrm{C}_{3} \mathrm{H}_{8}\) or \(\mathrm{CH}_{3} \mathrm{OCH}_{3}\), (c) \(\mathrm{HOOH}\) or (a) \(\mathrm{C}_{6} \mathrm{H}_{14}\) or \(\mathrm{C}_{8} \mathrm{H}_{18}\) \(\mathrm{HSSH}\), (d) \(\mathrm{NH}_{2} \mathrm{NH}_{2}\) or \(\mathrm{CH}_{3} \mathrm{CH}_{3}\)

Look up and compare the normal boiling points and normal melting points of \(\mathrm{H}_{2} \mathrm{O}\) and \(\mathrm{H}_{2} \mathrm{~S}\). (a) Based on these physical properties, which substance has stronger intermolecular forces? What kind of intermolecular forces exist for each molecule? (b) Predict whether solid \(\mathrm{H}_{2} \mathrm{~S}\) is more or less dense than liquid \(\mathrm{H}_{2} \mathrm{~S}\). How does this compare to \(\mathrm{H}_{2} \mathrm{O} ?\) Explain. (c) Water has an unusually high specific heat. Is this related to its intermolecular forces? Explain.

The following quote about ammonia \(\left(\mathrm{NH}_{3}\right)\) is from a textbook of inorganic chemistry: "It is estimated that \(26 \%\) of the hydrogen bonding in \(\mathrm{NH}_{3}\) breaks down on melting, \(7 \%\) on warming from the melting to the boiling point, and the final \(67 \%\) on transfer to the gas phase at the boiling point." From the standpoint of the kinetic energy of the molecules, explain (a) why there is a decrease of hydrogen-bonding energy on melting and (b) why most of the loss in hydrogen bonding occurs in the transition from the liquid to the vapor state.

(a) Explain why surface tension and viscosity decrease with increasing temperature. (b) Why do substances with high surface tensions also tend to have high viscosities?

(a) Distinguish between adhesive forces and cohesive forces. (b) What adhesive and cohesive forces are involved when a paper towel absorbs water? (c) Explain the cause for the U-shaped meniscus formed when water is in a glass tube.

Explain the following observations: (a) The surface tension of \(\mathrm{CHBr}_{3}\) is greater than that of \(\mathrm{CHCl}_{3}\). (b) As tem- perature increases, oil flows faster through a narrow tube. (c) Raindrops that collect on a waxed automobile hood take on a nearly spherical shape. (d) Oil droplets that collect on a waxed automobile hood take on a flat shape.

Hydrazine \(\left(\mathrm{H}_{2} \mathrm{NNH}_{2}\right)\), hydrogen peroxide (HOOH), and water \(\left(\mathrm{H}_{2} \mathrm{O}\right)\) all have exceptionally high surface tensions compared with other substances of comparable molecular weights. (a) Draw the Lewis structures for these three compounds. (b) What structural property do these substances have in common, and how might that account for the high surface tensions?

Name the phase transition in each of the following situations, and indicate whether it is exothermic or endothermic: (a) When ice is heated, it turns to water. (b) Wet clothes dry on a warm summer day. (c) Frost appears on a window on a cold winter day. (d) Droplets of water appear on a cold glass of beer.

Name the phase transition in each of the following situations, and indicate whether it is exothermic or endothermic: (a) Bromine vapor turns to bromine liquid as it is cooled. (b) Crystals of iodine disappear from an evaporating dish as they stand in a fume hood. (c) Rubbing alcohol in an open container slowly disappears. (d) Molten lava from a volcano turns into solid rock.

Explain why the heat of fusion of any substance is generally lower than its heat of vaporization.

Ethyl chloride \(\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}\right)\) boils at \(12^{\circ} \mathrm{C}\). When liquid \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}\) under pressure is sprayed on a room-temperature \(\left(25^{\circ} \mathrm{C}\right)\) surface in air, the surface is cooled considerably. (a) What does this observation tell us about the specific heat of \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}(g)\) as compared with \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Cl}(l) ?\) (b) As- sume that the heat lost by the surface is gained by ethyl chloride. What enthalpies must you consider if you were to calculate the final temperature of the surface?

For many years drinking water has been cooled in hot climates by evaporating it from the surfaces of canvas bags or porous clay pots. How many grams of water can be cooled from \(35^{\circ} \mathrm{C}\) to \(20^{\circ} \mathrm{C}\) by the evaporation of \(60 \mathrm{~g}\) of water? (The heat of vaporization of water in this temperature range is \(2.4 \mathrm{~kJ} / \mathrm{g}\). The specific heat of water is \(4.18 \mathrm{~J} / \mathrm{g}-\mathrm{K} .)\)

Compounds like \(\mathrm{CCl}_{2} \mathrm{~F}_{2}\) are known as chlorofluorocarbons, or CFCs. These compounds were once widely used as refrigerants but are now being replaced by compounds that are believed to be less harmful to the environment. The heat of vaporization of \(\mathrm{CCl}_{2} \mathrm{~F}_{2}\) is \(289 \mathrm{~J} / \mathrm{g}\). What mass of this substance must evaporate to freeze \(200 \mathrm{~g}\) of water initially at \(15^{\circ} \mathrm{C}\) ? (The heat of fusion of water is \(334 \mathrm{~J} / \mathrm{g} ;\) the specific heat of water is \(4.18 \mathrm{~J} / \mathrm{g}-\mathrm{K}\).)

Ethanol \(\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\right)\) melts at \(-114{ }^{\circ} \mathrm{C}\) and boils at \(78{ }^{\circ} \mathrm{C}\). Its density is \(0.789 \mathrm{~g} / \mathrm{mL}\). The enthalpy of fusion of ethanol is \(5.02 \mathrm{~kJ} / \mathrm{mol}\), and its enthalpy of vaporization is \(38.56 \mathrm{~kJ} / \mathrm{mol}\). The specific heats of solid and liquid ethanol are \(0.97 \mathrm{~J} / \mathrm{g}-\mathrm{K}\) and \(2.3 \mathrm{~J} / \mathrm{g}-\mathrm{K}\), respectively. (a) How much heat is required to convert \(25.0 \mathrm{~g}\) of ethanol at \(25^{\circ} \mathrm{C}\) to the vapor phase at \(78^{\circ} \mathrm{C} ?\) (b) How much heat is required to convert \(5.00 \mathrm{~L}\) of ethanol at \(-140^{\circ} \mathrm{C}\) to the vapor phase at \(78^{\circ} \mathrm{C} ?\)

The fluorocarbon compound \(\mathrm{C}_{2} \mathrm{Cl}_{3} \mathrm{~F}_{3}\) has a normal boiling point of \(47.6^{\circ} \mathrm{C}\). The specific heats of \(\mathrm{C}_{2} \mathrm{Cl}_{3} \mathrm{~F}_{3}(l)\) and \(\mathrm{C}_{2} \mathrm{Cl}_{3} \mathrm{~F}_{3}(g)\) are \(0.91 \mathrm{~J} / \mathrm{g}-\mathrm{K}\) and \(0.67 \mathrm{~J} / \mathrm{g}-\mathrm{K}\), respectively. The heat of vaporization for the compound is \(27.49 \mathrm{~kJ} / \mathrm{mol}\). Calculate the heat required to convert \(50.0 \mathrm{~g}\) of \(\mathrm{C}_{2} \mathrm{Cl}_{3} \mathrm{~F}_{3}\) from a liquid at \(10.00{ }^{\circ} \mathrm{C}\) to a gas at \(85.00{ }^{\circ} \mathrm{C}\).

Explain how each of the following affects the vapor pressure of a liquid: (a) volume of the liquid, (b) surface area, (c) intermolecular attractive forces, (d) temperature, (e) density of the liquid.

(a) Place the following substances in order of increasing volatility: \(\mathrm{CH}_{4}, \mathrm{CBr}_{4}, \mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CH}_{3} \mathrm{Cl}, \mathrm{CHBr}_{3}\), and \(\mathrm{CH}_{2} \mathrm{Br}_{2}\). Explain. (b) How do the boiling points vary through this series?

True or false: (a) \(\mathrm{CBr}_{4}\) is more volatile than \(\mathrm{CCl}_{4}\). (b) \(\mathrm{CBr}_{4}\) has a higher boiling point than \(\mathrm{CCl}_{4}\). (c) \(\mathrm{CBr}_{4}\) has weaker intermolecular forces than \(\mathrm{CCl}_{4}\). (d) \(\mathrm{CBr}_{4}\) has a higher vapor pressure at the same temperature than \(\mathrm{CCl}_{4}\)

(a) Two pans of water are on different burners of a stove. One pan of water is boiling vigorously, while the other is boiling gently. What can be said about the temperature of the water in the two pans? (b) A large container of water and a small one are at the same temperature. What can be said about the relative vapor pressures of the water in the two containers?

Explain the following observations: (a) Water evaporates more quickly on a hot, dry day than on a hot, humid day. (b) It takes longer to boil water for tea at high altitudes than at lower altitudes.

Appendix B lists the vapor pressure of water at various external pressures. (a) Plot the data in Appendix \(B\), vapor pressure (torr) vs. temperature \(\left({ }^{\circ} \mathrm{C}\right)\). From your plot, estimate the vapor pressure of water at body temperature, \(37^{\circ} \mathrm{C}\). (b) Explain the significance of the data point at \(760.0\) torr, \(100^{\circ} \mathrm{C}\) (c) A city at an altitude of \(5000 \mathrm{ft}\) above sea level has a barometric pressure of 633 torr. To what temperature would you have to heat water to boil it in this city? (d) A city at an altitude of \(500 \mathrm{ft}\) below sea level would have a barometric pressure of 774 torr. To what temperature would you have to heat water to boil it in this city? (e) For the two cities in parts (c) and (d), compare the average kinetic energies of the water molecules at their boiling points. Are the kinetic energies the same or different? Explain.

(a) What is the significance of the critical point in a phase diagram? (b) Why does the line that separates the gas and liquid phases end at the critical point?

(a) What is the significance of the triple point in a phase diagram? (b) Could you measure the triple point of water by measuring the temperature in a vessel in which water vapor, liquid water, and ice are in equilibrium under one atmosphere of air? Explain.

Sketch a generic phase diagram for a substance that has a more dense solid phase than a liquid phase. Label all regions, lines, and points.

The normal melting and boiling points of \(\mathrm{O}_{2}\) are \(-218^{\circ} \mathrm{C}\) and \(-183{ }^{\circ} \mathrm{C}\) respectively. Its triple point is at \(-219^{\circ} \mathrm{C}\) and \(1.14\) torr, and its critical point is at \(-119^{\circ} \mathrm{C}\) and \(49.8\) atm. (a) Sketch the phase diagram for \(\mathrm{O}_{2}\), showing the four points given and indicating the area in which each phase is stable. (b) Will \(\mathrm{O}_{2}(s)\) float on \(\mathrm{O}_{2}(t) ?\) Explain. (c) As it is heated, will solid \(\mathrm{O}_{2}\) sublime or melt under a pressure of 1 atm?

(a) Draw a picture that represents a crystalline solid at the atomic level. (b) Now draw a picture that represents an amorphous solid at the atomic level.

Amorphous silica has a density of about \(2.2 \mathrm{~g} / \mathrm{cm}^{3}\), whereas the density of crystalline quartz is \(2.65 \mathrm{~g} / \mathrm{cm}^{3}\) Account for this difference in densities.

An element crystallizes in a body-centered cubic lattice. The edge of the unit cell is \(2.86 \AA\), and the density of the crystal is \(792 \mathrm{~g} / \mathrm{cm}^{3}\), Calculate the atomic weight of the element.

\(\mathrm{KCl}\) has the same structure as \(\mathrm{NaCl}\). The length of the unit cell is \(628 \mathrm{pm}\). The density of \(\mathrm{KCl}\) is \(1.984 \mathrm{~g} / \mathrm{cm}^{3}\), and its formula mass is \(74.55\) amu. Using this information, calculate Avogadro's number.

A particular form of cinnabar (HgS) adopts the zinc blende structure, Figure \(11.42(b)\). The length of the unit cell side is \(5.852 \AA\). (a) Calculate the density of \(\mathrm{HgS}\) in this form. (b) The mineral tiemmanite (HgSe) also forms a solid phase with the zinc blende structure. The length of the unit cell side in this mineral is \(6.085 \AA\). What accounts for the larger unit cell length in tiemmanite? (c) Which of the two substances has the higher density? How do you account for the difference in densities?

It is possible to change the temperature and pressure of a vessel containing argon gas so that the gas solidifies. (a) What intermolecular forces exist between argon atoms? (b) Is the solid argon a "covalent network solid"? Why or why not?

(a) Silicon is the fundamental component of integrated circuits. Si has the same structure as diamond. Is \(\mathrm{Si}\) a molecular, metallic, ionic, or covalent-network solid? (b) Silica is \(\mathrm{SiO}_{2}\). What type of solid would you expect silica to form?

What kinds of attractive forces exist between particles in (a) molecular crystals, (b) covalent-network crystals, (c) ionic crystals, (d) metallic crystals?

Indicate the type of crystal (molecular, metallic, covalent-network, or ionic) each of the following would form upon solidification: (a) \(\mathrm{CaCO}_{3}\), (b) \(\mathrm{Pt}\), (c) \(\mathrm{ZrO}_{2}\) (melting point, \(2677^{\circ} \mathrm{C}\) ), (d) table sugar \(\left(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}\right)\) (e) benzene, (f) \(\mathrm{I}_{2}\).

Covalent bonding occurs in both molecular and covalent-network solids. Why do these two kinds of solids differ so greatly in their hardness and melting points?

Which type (or types) of crystalline solid is characterized by each of the following: (a) high mobility of electrons throughout the solid; (b) softness, relatively low melting point; (c) high melting point and poor electrical conductivity; (d) network of covalent bonds; (e) charged particles throughout the solid.