Chemistry The Molecular Science

John W. Moore, Conrad L. Stanitski

Chapter 9

Liquids, Solids, and Materials - all with Video Answers

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Chapter Questions

Problem 1

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

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Problem 2

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?

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Problem 3

Explain how the equilibrium vapor pressure of a liquid might be measured.

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Problem 4

Define boiling point and normal boiling point.

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Problem 5

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

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Problem 6

Define sublimation.

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Problem 7

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

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Problem 8

Define the unit cell of a crystal.

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Problem 9

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.

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Problem 10

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.

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Problem 11

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) $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OCH}_{3}$

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Problem 12

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) $\mathrm{SO}_{2}$

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Problem 13

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

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Problem 14

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?

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Problem 15

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.

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Problem 16

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 $25^{\circ} \mathrm{C}$.

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Problem 17

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.

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Problem 18

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 $40.7 \mathrm{~kJ} / \mathrm{mol}$.

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Problem 19

Use the concepts of noncovalent intermolecular forces to explain why vaporization is endothermic.

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Problem 20

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

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Problem 21

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

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Problem 22

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 $7.4{ }^{\circ} \mathrm{C}$

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Problem 23

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.

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Problem 24

The highest mountain in the Western Hemisphere is $\mathrm{Mt}$. Aconcagua, in the central Andes of Argentina $(22,834 \mathrm{ft})$ Assume that atmospheric pressure decreases at a rate of 3.5 millibar every 100 feet. (a) Estimate the atmospheric pressure at the summit of Mt. Aconcagua. (b) Calculate the temperature at which water would boil at the summit.

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Problem 25

The lowest sea-level barometric pressure ever recorded was 25.69 in mercury, recorded in a typhoon in the South Pacific. Suppose you were in this typhoon and, to calm yourself, boiled water to make yourself a cup of tea. At what temperature would the water boil? Use Figure 9.4 and remember that $1 \mathrm{~atm}$ is $760 \mathrm{mmHg}$ and 1 in = $2.54 \mathrm{~cm}$.

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Problem 26

The vapor pressure curves for four substances are shown in the plot. Which one of these four substances has the greatest intermolecular attractive forces at $25^{\circ} \mathrm{C} ?$ Explain your answer.

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Problem 27

A liquid has a $\Delta_{\text {vap }} H$ of $38.7 \mathrm{~kJ} / \mathrm{mol}$ and a boiling point of $110^{\circ} \mathrm{C}$ at 1 atm pressure. Calculate the vapor pressure of the liquid at $97^{\circ} \mathrm{C}$.

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Problem 28

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 $130^{\circ} \mathrm{C}$.

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Problem 29

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.

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Problem 30

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 $80.0^{\circ} \mathrm{C}$

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Problem 31

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

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Problem 32

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

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Problem 33

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

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Problem 34

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.

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Problem 35

Calculate the total heat energy transfer required to change 0.50 mol ice at $-5^{\circ} \mathrm{C}$ to 0.50 mol steam at $100^{\circ} \mathrm{C}$. The solid ice and liquid water have heat capacities of $2.06 \mathrm{~J} \mathrm{~g}^{-1}{ }^{\circ} \mathrm{C}^{-1}$ and $4.184 \mathrm{~J} \mathrm{~g}^{-1}{ }^{\circ} \mathrm{C}^{-1}$, respectively. The fusion enthalpy for solid ice is $6.02 \mathrm{~kJ} / \mathrm{mol}$ and the vaporization enthalpy of liquid water is $40.7 \mathrm{~kJ} / \mathrm{mol}$

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Problem 36

Calculate the heat energy transfer needed to melt a $36.00-\mathrm{g}$ ice cube that is initially at $-10^{\circ} \mathrm{C}$ and bring it to room temperature $\left(20^{\circ} \mathrm{C}\right)$. (See Question 35 for data.)

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Problem 37

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 $4.184 \mathrm{~J} \mathrm{~g}^{-1}{ }^{\circ} \mathrm{C}^{-1}$.

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Problem 38

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 $) ?$

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Problem 39

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.

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Problem 40

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) $\mathrm{CH}_{3} \mathrm{OH}$

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Problem 41

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) $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}$

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Problem 42

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?

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Problem 43

In this phase diagram, make these identifications:
(a) What phase is present in region A? Region B? Region C?
(b) What phases are in equilibrium at point 1 ? Point $2 ?$ Point 3 ? Point $5 ?$

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Problem 44

Consult the phase diagram of $\mathrm{CO}_{2}$ in Figure 9.14 . What phase or phases are present under these conditions:
(a) $T=-70^{\circ} \mathrm{C}$ and $P=1.0 \mathrm{~atm} .$
(b) $T=-40^{\circ} \mathrm{C}$ and $P=15.5 \mathrm{~atm} .$
(c) $T=-80^{\circ} \mathrm{C}$ and $P=4.7 \mathrm{~atm} .$

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Problem 45

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 $94 \mathrm{~cm}^{3}$

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Problem 46

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

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Problem 47

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

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Problem 48

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

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Problem 49

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

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Problem 50

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

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Problem 51

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 $\mathrm{NaCl}$

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Problem 52

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?

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Problem 53

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.

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Problem 54

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

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Problem 55

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

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Problem 56

Water can participate in hydrogen-bonding with other water molecules. In liquid water, how many hydrogen bonds does each water molecule engage in? What threedimensional shape do these bonds assume?

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Problem 57

Each diagram given represents an array of like atoms that would extend indefinitely in two dimensions. Draw a two-dimensional unit cell for each array. How many atoms are in each unit cell?

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Problem 58

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

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Problem 59

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.

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Problem 60

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

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Problem 61

Using the $\mathrm{NaCl}$ structure shown in Figure 9.24 , how many unit cells share each of the $\mathrm{Na}^{+}$ ions in the front face of the unit cell? How many unit cells share each of the $\mathrm{Cl}^{-}$ ions in this face?

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Problem 62

The ionic radii of $\mathrm{Cs}^{+}$ and $\mathrm{Cl}^{-}$ are 181 and $167 \mathrm{pm}$, respectively. What is the length of the body diagonal in the $\mathrm{CsCl}$ unit cell? What is the length of the side of this unit cell? (CsCl has the same unit cell as CsI, shown in Figure $9.23 .$ )

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Problem 63

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

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Problem 64

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.

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Problem 65

Solid lithium has a body-centered cubic unit cell with the length of the edge of $351 \mathrm{pm}$ at $20{ }^{\circ} \mathrm{C}$. Calculate the density of lithium at this temperature.

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Problem 66

Tungsten has a body-centered cubic unit cell and an atomic radius of $141 \mathrm{pm} .$ Calculate the density of solid tungsten.

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Problem 67

Explain why diamond is more dense than graphite.

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Problem 68

Explain how two-dimensional and three-dimensional network solids differ on the nanoscale.

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Problem 69

Determine, by looking up data in a reference such as the Handbook of Chemistry and Physics, whether the examples of network solids given in the text are soluble in water or other common solvents. Explain your answer in terms of the chemical bonding in network solids.

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Problem 70

Explain why diamond is an electrical insulator and graphite is an electrical conductor.

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Problem 71

For a clear diffraction pattern to be seen from a regularly spaced lattice, the radiation falling on the lattice must have a wavelength less than the lattice spacing. From the unit cell size of the $\mathrm{NaCl}$ crystal, estimate the maximum wavelength of the radiation that would be diffracted by this crystal. Calculate the frequency of the radiation and the energy associated with (a) one photon and (b) one mole of photons of the radiation. In what region of the spectrum is this radiation?

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Problem 72

Taking the middle of the visible spectrum to be green light with a wavelength of $550 \mathrm{nm}$, calculate how many aluminum atoms (radius $=143 \mathrm{pm}$ ) touching their neighbors would make a straight line $550 \mathrm{nm}$ long. Using this result, explain why an optical microscope using visible radiation will never be able to detect an individual aluminum atom (or any other atom, for that matter).

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Problem 73

The second-order Bragg reflection $(n=2)$ from a copper crystal for X-rays with a wavelength of $166 \mathrm{pm}$ is $27.35^{\circ}$. Calculate the spacing between the planes of copper atoms.

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Problem 74

The first-order Bragg reflection $(n=1)$ from a $\mathrm{NaCl}$ crystal with a spacing of $282 \mathrm{pm}$ is seen at $23.0^{\circ} .$ Calculate the wavelength of the $\mathrm{X}$ -ray radiation used.

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Problem 75

List the four major classes of materials and give one example of each.

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Problem 76

Identify three items you use daily that have been developed recently using the principles of materials science.

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Problem 77

What is the principal difference between the orbitals that electrons occupy in individual, isolated atoms and the orbitals they occupy in solid metals?

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Problem 78

In terms of band theory, what is the difference between a conductor and an insulator? Between a conductor and a semiconductor?

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Problem 79

Name three properties of metals, and explain them by using a theory of metallic bonding.

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Problem 80

Which substance has the greatest electrical conductivity? The smallest electrical conductivity? Explain your choice briefly.
(a) Si
(b) Ge
(c) $\mathrm{Ag}$
(d) $\mathrm{P}_{4}$

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Problem 81

Which substance has the greatest electrical conductivity? The smallest electrical conductivity? Explain your choices briefly.
(a) $\mathrm{RbCl}(\ell)$
(b) $\mathrm{NaBr}(\mathrm{s})$
(c) $\mathrm{Rb}$
(d) Diamond

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Problem 82

Define the term "superconductor." Give the chemical formulas of two kinds of superconductors and their associated transition temperatures.

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Problem 83

Extremely high-purity silicon is required to manufacture semiconductors such as the memory chips found in calculators and computers. If a silicon wafer is $99.99999999 \%$ pure, approximately how many atoms of some other element are present per gram of high-purity silicon?

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Problem 84

Explain why Group $3 \mathrm{~A}$ and Group $5 \mathrm{~A}$ elements are used to dope silicon to improve its semiconducting properties.

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Problem 85

Explain the difference between $n$ -type semiconductors and $p$ -type semiconductors.

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Problem 86

Define the term "amorphous."

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Problem 87

What makes a glass different from a crystalline solid such as $\mathrm{SiO}_{2}$ ? Under what conditions could $\mathrm{SiO}_{2}$ become glass-like?

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Problem 88

A typical cement contains, by weight, $65 \% \mathrm{CaO}, 20 \%$ $\mathrm{SiO}_{2}, 5 \% \mathrm{Al}_{2} \mathrm{O}_{3}, 6 \% \mathrm{Fe}_{2} \mathrm{O}_{3},$ and $4 \% \mathrm{MgO} .$ Determine the mass percent of each type of atom present. Then, determine an empirical formula of the material from the percent composition, setting the subscript of the least abundant element to 1.00 .

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Problem 89

Give two examples of (a) silicate ceramics, (b) oxide ceramics, and (c) nonoxide ceramics.

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Problem 90

Explain why, when you boil water in a pan, the water boils much faster when the pan has a lid on it than when it does not.

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Problem 91

Will a closed container of water at $70^{\circ} \mathrm{C}$ or an open container of water at the same temperature cool faster on a cold winter day? Explain why.

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Problem 92

Given these properties: Camphor: colorless needles; density $=0.900 \mathrm{~g} / \mathrm{cm}^{3}$ at $25^{\circ} \mathrm{C}$; sublimes at $204{ }^{\circ} \mathrm{C} ;$ insoluble in water; very soluble in ethanol or ether. Praseodymium chloride: blue-green needle crystals; density $=4.02 \mathrm{~g} / \mathrm{cm}^{3}$ at $25^{\circ} \mathrm{C}$; melting point $786^{\circ} \mathrm{C}$; boiling point $1700^{\circ} \mathrm{C}$; solubility $103.9 \mathrm{~g} / 100 \mathrm{~mL}$ cold water, very soluble in hot water.
(a) Is camphor an ionic or covalent compound? Explain
your answer.
(b) Is praseodymium chloride an ionic or covalent compound? Explain your answer.

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Problem 93

Xenon has a triple point of $0.81 \mathrm{~atm}$ and $-112{ }^{\circ} \mathrm{C}$ and a normal boiling point of $-108^{\circ} \mathrm{C}$. If the pressure exerted on a xenon sample is $1.75 \mathrm{~atm}$ and the temperature is $-105^{\circ} \mathrm{C}$, in what phase (physical state) does the xenon sample likely exist at these conditions? Explain your answer.

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Problem 94

Xenon has a triple point of $0.81 \mathrm{~atm}$ and $-112{ }^{\circ} \mathrm{C}$ and a normal boiling point of $-108^{\circ} \mathrm{C}$. If the pressure exerted on a xenon sample is $1.75 \mathrm{~atm}$ and the temperature is $-105^{\circ} \mathrm{C}$, in what phase (physical state) does the xenon sample likely exist at these conditions? Explain your answer.

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Problem 95

Use the vapor pressure curves for methyl ethyl ether, $\mathrm{CH}_{3} \mathrm{OCH}_{2} \mathrm{CH}_{3} ;$ carbon disulfide, $\mathrm{CS}_{2}$; and benzene, $\mathrm{C}_{6} \mathrm{H}_{6},$ to answer these questions.
(a) What is the vapor pressure of methyl ethyl ether at $0{ }^{\circ} \mathrm{C} ?$
(b) Which of these three liquids has the strongest intermolecular attractions?
(c) At what temperature does benzene have a vapor pressure of $600 \mathrm{mmHg}$ ?
(d) What are the normal boiling points of these three liquids?

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Problem 96

Organic compounds with structures based on benzene, $\mathrm{C}_{6} \mathrm{H}_{6},$ can be formed by substituting an atom or a group of atoms for one of the hydrogens. Such substituted benzenes have their own properties, different from benzene and from each other. Explain the order of experimental boiling points for these four compounds.
(a) $\mathrm{C}_{6} \mathrm{H}_{6}\left(80{ }^{\circ} \mathrm{C}\right)$
(b) $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{Cl}\left(131{ }^{\circ} \mathrm{C}\right)$
(c) $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{Br}\left(156^{\circ} \mathrm{C}\right)$
(d) $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}\left(182{ }^{\circ} \mathrm{C}\right)$

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Numerade Educator

Problem 97

The chlorofluorocarbon $\mathrm{CCl}_{2} \mathrm{~F}_{2}$ was once used in air conditioners as the heat transfer fluid. Its normal boiling point is $-30^{\circ} \mathrm{C}$, and its vaporization enthalpy is $165 \mathrm{~J} \mathrm{~g}^{-1}$. The gas and the liquid have specific heat capacities of $0.61 \mathrm{~J} \mathrm{~g}^{-1}{ }^{\circ} \mathrm{C}^{-1}$ and $0.97 \mathrm{~J} \mathrm{~g}^{-1}{ }^{\circ} \mathrm{C}^{-1}$ respectively. Calculate the heat energy transfer when $10.0 \mathrm{~g} \mathrm{CCl}_{2} \mathrm{~F}_{2}$ is cooled from $40^{\circ} \mathrm{C}$ to $-40^{\circ} \mathrm{C}$.

Pronoy Sinha

Numerade Educator

Problem 98

Liquid ammonia, $\mathrm{NH}_{3}(\ell)$, was used as a refrigerant fluid before the discovery of the chlorofluorocarbons and is still widely used today. Its normal boiling point is $-33.4^{\circ} \mathrm{C},$ and its vaporization enthalpy is $23.5 \mathrm{~kJ} / \mathrm{mol}$. The gas and liquid have specific heat capacities of 2.2 $\mathrm{J} \mathrm{g}^{-1} \mathrm{~K}^{-1}$ and $4.7 \mathrm{~J} \mathrm{~g}^{-1} \mathrm{~K}^{-1}$, respectively. Calculate the heat energy transfer required to raise the temperature of $10.0 \mathrm{~kg}$ liquid ammonia from $-50.0^{\circ} \mathrm{C}$ to $0.0{ }^{\circ} \mathrm{C}$.

Aadit Sharma

Numerade Educator

Problem 99

Sulfur dioxide, $\mathrm{SO}_{2}$, is found in polluted air. What types of forces are responsible for binding $\mathrm{SO}_{2}$, molecules to one another in the solid or liquid phase?

Pronoy Sinha

Numerade Educator

Problem 100

Using the information below, place the compounds listed in order of increasing intermolecular attractions. For each substance list all types of intermolecular forces that are important.
$$
\begin{array}{lc}
\hline \text { Compound } & \text { Normal Boiling Point }\left({ }^{\circ} \mathrm{C}\right) \\
\hline \mathrm{SO}_{2} & -10 \\
\mathrm{NH}_{3} & -33.4 \\
\mathrm{CH}_{4} & -161.5 \\
\mathrm{H}_{2} \mathrm{O} & 100 \\
\hline
\end{array}
$$

Pronoy Sinha

Numerade Educator

Problem 101

Metallic gold is very malleable; that is, it can be hammered into very thin sheets, which are sometimes called gold leaf. For example, a 1.0 -g sample of metallic gold can be hammered into a sheet with an area of $1.0 \mathrm{~m}^{2}$. The density of gold is $19.3 \mathrm{~g} / \mathrm{cm}^{3}$ and the radius of a gold atom is $144 \mathrm{pm}$. Calculate how many atoms thick such a gold sheet would be.

Pronoy Sinha

Numerade Educator

Problem 102

Explain why, in general, the vaporization enthalpy of a liquid is much greater than the fusion enthalpy of its solid.

Pronoy Sinha

Numerade Educator

Problem 103

Consider this information regarding two compounds (common names are used). Orpiment: yellow solid; density $=3.49 \mathrm{~g} / \mathrm{cm}^{3}$ at $25^{\circ} \mathrm{C} ;$ melting point $=573 \mathrm{~K}$; slightly soluble in hot water; soluble in basic solution. Zeaxanthin: orange-red solid; density $=0.93 \mathrm{~g} / \mathrm{cm}^{3}$ at $25^{\circ} \mathrm{C} ;$ melting point $=489 \mathrm{~K} ;$ insoluble in water, soluble in benzene.
(a) Is orpiment an ionic or molecular compound? Explain your answer.
(b) Is zeaxanthin an ionic or molecular compound? Explain your answer.

Pronoy Sinha

Numerade Educator

Problem 104

Consider this information regarding two compounds. Thallium azide: yellow crystalline solid; melting point = $330^{\circ} \mathrm{C}$; slightly soluble in water, more soluble in hot water; insoluble in ethanol or diethyl ether. Camphene:
colorless, cubic crystals; melting point $=51{ }^{\circ} \mathrm{C}$; boiling point $=159{ }^{\circ} \mathrm{C} ;$ insoluble in water; moderately soluble in ethanol; soluble in diethyl ether.
(a) Is camphene an ionic or molecular compound? Explain your answer.
(b) Is thallium azide an ionic or molecular compound? Explain your answer.

Pronoy Sinha

Numerade Educator

Problem 105

The phase diagram for red phosphorus is shown nearby.
(a) Label the areas for the pure solid, liquid, and gas phases.
(b) Label a point on the diagram at which red phosphorus liquid and vapor are in equilibrium.
(c) Explain why solid red phosphorus cannot be melted in a container open to the atmosphere.
(d) Identify all the phase changes that occur sequentially when conditions change from Point $\mathrm{B}$ to Point $\mathrm{A}$.

Aadit Sharma

Numerade Educator

Problem 106

Consider liquid water in equilibrium with its vapor at $100{ }^{\circ} \mathrm{C}$. Estimate the number of water molecules per $\mathrm{cm}^{3}$ in (a) the liquid (density $=0.958 \mathrm{~g} / \mathrm{cm}^{3}$ at $\left.100{ }^{\circ} \mathrm{C}\right)$; (b) the vapor.

Pronoy Sinha

Numerade Educator

Problem 107

If you get boiling water at $100^{\circ} \mathrm{C}$ on your skin, it burns. If you get $100^{\circ} \mathrm{C}$ steam on your skin, it burns much more severely. Explain why this is so.

Aadit Sharma

Numerade Educator

Problem 108

If water at room temperature is placed in a flask that is connected to a vacuum pump and the vacuum pump then lowers the pressure in the flask, we observe that the volume of the water decreases and the remaining water turns into ice. Explain what has happened.

Aadit Sharma

Numerade Educator

Problem 109

The normal boiling point of $\mathrm{SO}_{2}$ is $263.1 \mathrm{~K}$ and that of $\mathrm{NH}_{3}$ is $239.7 \mathrm{~K}$. At $-40^{\circ} \mathrm{C}$, would you predict that ammonia has a vapor pressure greater than, less than, or equal to that of sulfur dioxide? Explain.

Pronoy Sinha

Numerade Educator

Problem 110

Butane is a gas at room temperature; however, if you look closely at a butane lighter you see it contains liquid butane. Explain how it is possible to have liquid butane present.

Pronoy Sinha

Numerade Educator

Problem 111

While camping with a friend in the Rocky Mountains, you decide to cook macaroni for dinner. Your friend says the macaroni will cook faster in the Rockies because the lower atmospheric pressure will cause the water to boil at a lower temperature. Do you agree with your friend? Explain your reasoning.

Pronoy Sinha

Numerade Educator

Problem 112

Examine the nanoscale diagrams and the phase diagram for Question 112. Match each particulate diagram (1 through 7) to its corresponding point (A through H) on the phase diagram.

Aadit Sharma

Numerade Educator

Problem 113

Consider the phase diagram and heating-curve graphs for Question $113 .$ Draw corresponding heating curves for $T_{1}$ to $T_{2}$ at pressures $P_{1}$ and $P_{2} .$ Label each phase and phase change on your heating curves.

Aadit Sharma

Numerade Educator

Problem 114

Consider three boxes of equal volume. One is filled with tennis balls, another with golf balls, and the third with marbles. If a closest-packing arrangement is used in each box, which one has the most occupied space? Which one has the least occupied space? (Disregard the difference in filling space at the walls, bottom, and top of the box.)

Aadit Sharma

Numerade Educator

Problem 115

When spherical atoms are positioned in a unit cell, they don't occupy all of the space available in the unit cell. Some types of unit cells have more efficient packing than other types have. Consider two metals: one with a bodycentered cubic unit cell; the other with a face-centered cubic unit cell. Calculate the volume of the atoms in each of these unit cells in comparison to the volume of the unit cell itself. From these data, calculate the fraction of space in each of the unit cells that is occupied by its atoms. The volume of a sphere is $\frac{4}{3} \pi r^{3}$ where $r$ is the radius of the sphere.

Nicole Smina

Numerade Educator

Problem 116

Rhombic and monoclinic are two allotropic forms of sulfur that can exist as separate phases. The phase diagram for sulfur is shown nearby. Because of the wide range of pressure, the pressure scale is logarithmic.
(a) How many solid phases are there? Explain.
(b) How many triple points are there? Explain.
(c) In what phase(s) does sulfur exist(s) at 1 atm and $80^{\circ} \mathrm{C}$ ? At $125^{\circ} \mathrm{C}$ ? Explain your answers.
(d) What phases is sulfur in at $151^{\circ} \mathrm{C} ?$ Explain your answer.
(e) Based on this phase diagram, under what temperature and pressure conditions will sulfur sublime?
(f) What is the most stable phase at 1 atm and $100^{\circ} \mathrm{C}$ ?
(g) Identify in sequence the phases present at $10^{-4}$ atm as the temperature changes from $50^{\circ} \mathrm{C}$ to $200^{\circ} \mathrm{C}$.
(h) Determine the normal melting point of sulfur.

Matthew Confer

Numerade Educator

Problem 117

The phase diagram for water over a relative narrow pressure and temperature range is given in Figure 9.19 . A phase diagram over a considerably wider range of temperature and pressure (kbar) is given nearby. This phase diagram illustrates the polymorphism of ice, the existence of a solid in more than one form. In this case, Roman numerals are used to designate each polymorphic form. For example, Ice I, ordinary ice, is the form that exists under ordinary pressures. The other forms exist only at higher pressures, in some cases extremely high pressure such as Ice VII and Ice VIII.
(a) Using the phase diagram, give the approximate $P$ and $T$ conditions at the triple point for Ice III, Ice $\mathrm{V}$, and liquid water
(b) Determine the approximate temperature and pressure for the triple point for Ices VI, VII, and VIII.
(c) What is anomalously different about the fusion curves for Ice VI and Ice VII compared to that of Ice I?
(d) What phases exist at 8 kbar and $20{ }^{\circ} \mathrm{C}$ ?
(e) At a constant temperature of $-10^{\circ} \mathrm{C}$, start at $3 \mathrm{kbar}$ and increase the pressure to 7 kbar. Identify all the phase changes that occur sequentially as these conditions change.
(f) Explain why there is no triple point for the combination of Ice VII, Ice VIII, and liquid water.

Christina Sanchez

Christina Sanchez

Numerade Educator

Problem 118

Barium sulfide(s) has the $\mathrm{NaCl}$ structure and a density of $4.25 \mathrm{~g} / \mathrm{cm}^{3} .$ Calculate the interionic distance and compare this value with the sum of the ionic radii $\left(\mathrm{Ba}^{2+}=\right.$ $\left.149 \mathrm{pm} ; \mathrm{S}^{2-}=170 \mathrm{pm}\right)$

Vipin Singh

Numerade Educator

Problem 119

The Ne atom and the molecules $\mathrm{HF}, \mathrm{H}_{2} \mathrm{O}, \mathrm{NH}_{3},$ and $\mathrm{CH}_{4}$ all have the same number of electrons. In a thought experiment, you can make HF from Ne by removing a single proton a short distance from the nucleus and having the electrons follow the new arrangement of nuclei so as to make a new chemical bond. You can do the same for each of the other molecules. (Of course, none of these thought experiments can actually be done because of the enormous energies required to remove protons from nuclei.) For all of these substances, make a plot of (a) the boiling point in kelvins versus the number of hydrogen atoms, and (b) the molar vaporization enthalpy versus the number of hydrogen atoms. Explain any trend that you see in terms of intermolecular forces.

Linda Winkler

Numerade Educator

Problem 120

Metallic lithium has a body-centered cubic structure, and its unit cell is 351 pm along an edge. Lithium iodide has the same crystal lattice structure as sodium chloride (Figure 9.24 ). The cubic unit cell is 600 pm along an edge.
(a) Assume that the metal atoms in lithium touch along the body diagonal of the body-centered cubic unit cell, and estimate the radius of a lithium atom.
(b) Assume that in lithium iodide the $\mathrm{I}^{-}$ ions touch along the face diagonal of the cubic unit cell and that the $\mathrm{Li}^{+}$ and $\mathrm{I}^{-}$ ions touch along the edge of the cube; calculate the radius of an $\mathrm{I}^{-}$ ion and of an $\mathrm{Li}^{+}$ ion.
(c) Compare your results in parts (a) and
(b) for the radius of a lithium atom and a lithium ion. Are your results reasonable? If not, how could you account for the unexpected result? Could any of the assumptions that were made be in error? Explain.

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 121

Potassium chloride and rubidium chloride both have the sodium chloride structure (Figure 9.24 ). X-ray diffraction experiments indicate that their cubic unit cell dimensions are $629 \mathrm{pm}$ and $658 \mathrm{pm}$, respectively.
(i) One $m o l \mathrm{KCl}$ and $1 \mathrm{~mol} \mathrm{RbCl}$ are ground together to a very fine powder in a mortar and pestle, and the X-ray diffraction pattern of the pulverized solid is measured. Two patterns are observed, each corresponding to a cubic unit cell-one with an edge length of $629 \mathrm{pm}$ and one with an edge length of $658 \mathrm{pm}$. Call this Sample 1 .
(ii) One $\mathrm{mol} \mathrm{KCl}$ and $1 \mathrm{~mol} \mathrm{RbCl}$ are heated until the entire mixture is molten and then cooled to room temperature. A single X-ray diffraction pattern indicates a cubic unit cell with an edge length of roughly $640 \mathrm{pm}$. Call this Sample 2 .
(a) Suppose that Samples 1 and 2 were analyzed for their chloride content. What fraction of each sample is chloride? Could the samples be distinguished by means of chemical analysis?
(b) Interpret the two X-ray diffraction results in terms of the structures of the crystal lattices of Samples 1 and 2 .
(c) What chemical formula should you write for Sample $1 ?$ For Sample $2 ?$
(d) Suppose that you dissolved $1.00 \mathrm{~g}$ Sample 1 in $100 \mathrm{~mL}$ water in a beaker and did the same with $1.00 \mathrm{~g}$ Sample
2. Which sample would conduct electricity better, or would both be the same? What ions would be present in each solution at what concentrations?

David Collins

Numerade Educator

Problem 122

In lithium chloride, the larger $\mathrm{Cl}^{-}$ ions form a facecentered cubic unit cell lattice with the $\mathrm{Li}^{+}$ ions fitting between the $\mathrm{Cl}^{-}$ ions at the center of each edge of the unit cell. Sodium chloride has a similar structure. As shown in Figure 9.24 , the $\mathrm{Na}^{+}$ and $\mathrm{Cl}^{-}$ ions in sodium chloride are in a face-centered cubic unit cell in which the $\mathrm{Na}^{+}$ and $\mathrm{Cl}^{-}$ ions touch along each cell edge. The $\mathrm{Cl}^{-}$ ions are also at the centers of each face. Whether cations can fit into a particular unit cell lattice of anions to create a stable structure depends on the relative sizes of the cations and anions.
(a) Use ionic radii to calculate the $r_{\text {cation }} / r_{\text {anion }}$ ratio at which a cation just fits into a face-centered unit cell.
(b) How closely does sodium chloride meet this criterion?
(c) Comment on how well LiCl meets this criterion. Explain the difference between $\mathrm{LiCl}$ and $\mathrm{NaCl}$ in this regard.

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 123

Titanium metal crystallizes in a body-centered unit cell with a cell edge length of $330.6 \mathrm{pm}$. The density of titanium metal at $25.0^{\circ} \mathrm{C}$ is $4.401 \mathrm{~g} / \mathrm{cm}^{3} .$ Use these data to calculate the Avogadro constant.

Pronoy Sinha

Numerade Educator