Chapter 9

Name three properties of solids that are different from those of liquids. Explain the differences for each.

What causes surface tension in liquids? Name a substance that has a very high surface tension. What kinds of intermolecular forces account for the high value?

Define boiling point and normal boiling point.

Define the crystallization enthalpy of a substance. How is it related to the substance's fusion enthalpy?

Define sublimation.

Which processes are endothermic? (a) Condensation (b) Melting (c) Evaporation (d) Sublimation (e) Deposition (f) Freezing

Define the unit cell of a crystal.

Assuming the same substance could form crystals with its atoms or ions in either primitive cubic packing or hexagonal closest packing, which form would have the higher density? Explain.

How does conductivity vary with temperature for (a) a metallic conductor, (b) a nonconductor, (c) a semiconductor, and (d) a superconductor? In your answer, begin at high temperatures and come down to low temperatures.

Rank these substances in order of increasing noncovalent intermolecular attractions. For each substance, name the types of intermolecular attractions that occur. (a) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}\) (b) \(\mathrm{CH}_{3} \mathrm{CH}_{3}\) (c) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH}\) (d) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OCH}_{3}\)

Rank these substances in order of increasing noncovalent intermolecular attractions. For each substance, name the types of intermolecular attractions that occur. (a) \(\mathrm{Cl}_{2}\) (b) \(\mathrm{HF}\) (c) \(\mathrm{F}_{2}\) (d) \(\mathrm{SO}_{2}\)

Explain on the molecular scale the processes of condensation and vaporization.

After exercising on a hot summer day and working up a sweat, you often become cool when you stop. What is the molecular-level explanation of this phenomenon?

The chlorofluorocarbon \(\mathrm{CCl}_{3} \mathrm{~F}\) has a vaporization enthalpy of \(24.8 \mathrm{~kJ} / \mathrm{mol}\). Calculate the heat energy transfer required to vaporize \(1.00 \mathrm{~kg}\) of the compound.

The molar vaporization enthalpy of methanol is \(38.0 \mathrm{~kJ} / \mathrm{mol}\) at \(25^{\circ} \mathrm{C} .\) Calculate the heat energy transfer required to convert \(250 . \mathrm{mL}\) of the alcohol from liquid to vapor. The density of \(\mathrm{CH}_{3} \mathrm{OH}\) is \(0.787 \mathrm{~g} / \mathrm{mL}\) at \(25^{\circ} \mathrm{C}\).

Some camping stoves contain liquid butane, \(\mathrm{C}_{4} \mathrm{H}_{10} .\) They work only when the outside temperature is warm enough to allow the butane to have a reasonable vapor pressure (so they are not very good for camping in temperatures below about \(\left.0{ }^{\circ} \mathrm{C}\right)\). Assume the vaporization enthalpy of butane is \(22.44 \mathrm{~kJ} / \mathrm{mol}\) and the camp stove fuel tank contains \(190 . \mathrm{g}\) liquid \(\mathrm{C}_{4} \mathrm{H}_{10} .\) Calculate the heat energy transfer required to vaporize all of the butane.

Mercury is highly toxic. Although it is a liquid at room temperature, it has a high vapor pressure and a low vaporization enthalpy ( \(294 \mathrm{~J} / \mathrm{g}\) ). Calculate the heat energy transfer required to vaporize \(0.500 \mathrm{~mL}\) mercury at \(357^{\circ} \mathrm{C}\), its normal boiling point. The density of \(\operatorname{Hg}(\ell)\) is \(13.6 \mathrm{~g} / \mathrm{mL}\). Compare your result with the energy transfer needed to vaporize \(0.500 \mathrm{~mL}\) water. The molar vaporization enthalpy of \(\mathrm{H}_{2} \mathrm{O}\) is \(40.7 \mathrm{~kJ} / \mathrm{mol}\).

Give a molecular-level explanation of why the vapor pressure of a liquid increases with temperature.

Briefly explain the variations in the boiling points in this table. In your discussion be sure to mention the types of intermolecular forces involved. $$ \begin{array}{lc} \hline \text { Compound } & \text { Boiling Point }\left({ }^{\circ} \mathrm{C}\right) \\ \hline \mathrm{NH}_{3} & -33.4 \\ \mathrm{PH}_{3} & -87.5 \\ \mathrm{AsH}_{3} & -62.4 \\ \mathrm{SbH}_{3} & -17 \\ \hline \end{array} $$

Explain the observation that 1 -propanol, \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH},\) has a boiling point of \(97.2^{\circ} \mathrm{C},\) whereas a compound with the same empirical formula, ethyl methyl ether, \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OCH}_{3},\) boils at \(7.4{ }^{\circ} \mathrm{C}\)

Methanol, \(\mathrm{CH}_{3} \mathrm{OH}\), has a normal boiling point of \(64.7^{\circ} \mathrm{C}\) and a vapor pressure of \(100 \mathrm{mmHg}\) at \(21.2^{\circ} \mathrm{C}\). Formaldehyde, \(\mathrm{H}_{2} \mathrm{C}=\mathrm{O},\) has a normal boiling point of \(-19.5^{\circ} \mathrm{C}\) and a vapor pressure of \(100 \mathrm{mmHg}\) at \(-57.3^{\circ} \mathrm{C} .\) Explain why these two compounds have different boiling points and require different temperatures to achieve the same vapor pressure.

A liquid has a \(\Delta_{\text {vap }} H\) of \(44.0 \mathrm{~kJ} / \mathrm{mol}\) and a vapor pressure of \(370 \mathrm{mmHg}\) at \(90^{\circ} \mathrm{C}\). Calculate the vapor pressure of the liquid at \(130^{\circ} \mathrm{C}\).

The vapor pressure of ethanol, \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\), at \(50.0{ }^{\circ} \mathrm{C}\) is \(233 \mathrm{mmHg},\) and its normal boiling point at \(1 \mathrm{~atm}\) is \(78.3^{\circ} \mathrm{C} .\) Calculate the \(\Delta_{\mathrm{vap}} H\) of ethanol.

Calculate the \(\Delta_{\mathrm{vap}} H\) for a substance whose vapor pressure doubled when its temperature was raised from \(70.0^{\circ} \mathrm{C}\) to \(80.0^{\circ} \mathrm{C}\)

What does a low fusion enthalpy for a solid tell you about the solid (its bonding or type)?

What does a high melting point and a high fusion enthalpy tell you about a solid (its bonding or type)?

Which would you expect to have the higher fusion enthalpy, \(\mathrm{N}_{2}\) or \(\mathrm{I}_{2}\) ? Explain your choice.

The fusion enthalpy for \(\mathrm{H}_{2} \mathrm{O}\) is about 2.5 times larger than the fusion enthalpy for \(\mathrm{H}_{2} \mathrm{~S}\). What does this say about the relative strengths of the forces between the molecules in these two solids? Explain.

The chlorofluorocarbon \(\mathrm{CCl}_{2} \mathrm{~F}_{2}\) was once used as a refrigerant. Calculate what mass of this substance must evaporate to freeze \(2 \mathrm{~mol}\) water initially at \(20^{\circ} \mathrm{C}\). The vaporization enthalpy for \(\mathrm{CCl}_{2} \mathrm{~F}_{2}\) is \(289 \mathrm{~J} / \mathrm{g}\). The fusion enthalpy for solid ice is \(6.02 \mathrm{~kJ} / \mathrm{mol}\) and specific heat capacity for liquid water is \(4.184 \mathrm{~J} \mathrm{~g}^{-1}{\underline{\phantom{xx}}}^{\circ} \mathrm{C}^{-1}\).

The ions of \(\mathrm{NaF}\) and \(\mathrm{MgO}\) all have the same number of electrons, and the internuclear distances are about the same ( \(235 \mathrm{pm}\) and \(212 \mathrm{pm}\) ). Why, then, are the melting points of \(\mathrm{NaF}\) and \(\mathrm{MgO}\) so different \(\left(992{ }^{\circ} \mathrm{C}\right.\) and \(2825^{\circ} \mathrm{C},\) respectively \() ?\)

For the pair of compounds \(\mathrm{LiF}\) and \(\mathrm{CsI},\) tell which compound is expected to have the higher melting point and briefly explain why.

Which of these substances has the highest melting point? The lowest melting point? Explain your choices briefly. (a) \(\mathrm{LiBr}\) (b) \(\mathrm{CaO}\) (c) \(\mathrm{CO}\) (d) \(\mathrm{CH}_{3} \mathrm{OH}\)

Which of these substances has the highest melting point? The lowest melting point? Explain your choices briefly. (a) \(\mathrm{SiC}\) (b) I (c) \(\mathrm{Rb}\) (d) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}\)

During thunderstorms, very large hailstones can fall from the sky. To preserve some of these hailstones, you place them in the freezer compartment of your frost-free refrigerator. A friend, who is a chemistry student, tells you to put the hailstones in a tightly sealed plastic bag. Why?

At the critical point for carbon dioxide, the substance is very far from being an ideal gas. Prove this statement by calculating the density of an ideal gas in \(\mathrm{g} / \mathrm{cm}^{3}\) at the conditions of the critical point and comparing it with the experimental value. Compute the experimental value from the fact that a mole of \(\mathrm{CO}_{2}\) at its critical point occupies \(94 \mathrm{~cm}^{3}\)

Classify each of these solids as ionic, metallic, molecular, network, or amorphous. (a) \(\mathrm{KF}\) (b) \(\mathrm{I}_{2}\) (c) \(\mathrm{SiO}_{2}\) (d) \(\mathrm{BN}\)

Classify each of these solids as ionic, metallic, molecular, network, or amorphous. (a) Tetraphosphorus decaoxide (b) Brass (c) Ammonium phosphate (d) Graphite

On the basis of the description given, classify each of these solids as molecular, metallic, ionic, network, or amorphous, and explain your reasoning. (a) A brittle, yellow solid that melts at \(113^{\circ} \mathrm{C} ;\) neither the solid nor the liquid conducts electricity (b) A soft, silvery solid that melts at \(40^{\circ} \mathrm{C} ;\) both the solid and the liquid conduct electricity (c) A hard, colorless, crystalline solid that melts at \(1713^{\circ} \mathrm{C} ;\) neither the solid nor the liquid conducts electricity (d) A soft, slippery solid that melts at \(63^{\circ} \mathrm{C} ;\) neither the solid nor the liquid conducts electricity

On the basis of the description given, classify each of these solids as molecular, metallic, ionic, network, or amorphous, and explain your reasoning. (a) A soft, slippery solid that has no definite melting point but decomposes at temperatures above \(250{ }^{\circ} \mathrm{C} ;\) the solid does not conduct electricity (b) Violet crystals that melt at \(114{ }^{\circ} \mathrm{C}\) and whose vapor irritates the nose; neither the solid nor the liquid conducts electricity (c) Hard, colorless crystals that melt at \(2800^{\circ} \mathrm{C} ;\) the liquid conducts electricity, but the solid does not (d) A hard solid that melts at \(3410^{\circ} \mathrm{C}\); both the solid and the liquid conduct electricity

What type of solid exhibits each of these sets of properties? (a) Melts below \(100{ }^{\circ} \mathrm{C}\) and is insoluble in water (b) Conducts electricity only when melted (c) Insoluble in water and conducts electricity (d) Noncrystalline and melts over a wide temperature range

Describe how each of these materials would behave if it were deformed by a hammer strike. Explain why the materials behave as they do. (a) A metal, such as gold (b) A nonmetal, such as sulfur (c) An ionic compound, such as \(\mathrm{NaCl}\)

For most substances, the density of the solid phase is larger than for the liquid phase, but for water the reverse is true. What is the molecular-scale reason for this property of water? Why is this property important?

Explain how the changes of the density of water with temperature causes "turnover" in a lake in the spring and fall. Explain why the turnover is important.

The surface tension of a liquid decreases with increasing temperature. Using the idea of intermolecular attractions, explain why this is so.

The boiling point of water is relatively high for a compound of such low molar mass. Explain why this is so.

Name and draw the three cubic unit cells. Describe their similarities and differences.

Solid xenon forms crystals with a face-centered cubic unit cell that has an edge of \(620 \mathrm{pm} .\) Calculate the atomic radius of xenon.

Gold (atomic radius \(=144 \mathrm{pm}\) ) crystallizes in an \(\mathrm{fcc}\) unit cell. Calculate the length of a side of the cell.

Could \(\mathrm{CaCl}_{2}\) possibly have the \(\mathrm{NaCl}\) structure? Explain your answer briefly.

You know that thallium chloride, TlCl, crystallizes in either a primitive cubic or a face-centered cubic lattice of \(\mathrm{Cl}^{-}\) ions with \(\mathrm{Tl}^{+}\) ions in the holes. If the density of the solid is \(7.00 \mathrm{~g} / \mathrm{cm}^{3}\) and the edge of the unit cell is \(3.85 \times 10^{-8} \mathrm{~cm},\) determine the unit cell geometry.