Chemistry: The Molecular Science

John W. Moore, Conrad L. Stanitski, Peter C. Jurs

Chapter 11

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

List the postulates of the kinetic-molecular theory that apply to liquids.

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Problem 3

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 4

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

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Problem 5

Define boiling point and normal boiling point.

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Problem 6

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

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Problem 7

Define sublimation.

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Problem 8

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

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Problem 9

Define the unit cell of a crvstal

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Problem 10

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

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Problem 11

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 12

Predict which substance in Table $11.1$ has a surface tension most similar to that of each liquid:
(a) Ethylene glycol, $\mathrm{HOCH}_{2}-\mathrm{CH}_{2} \mathrm{OH}$
(b) Hexane, $\mathrm{C}_{6} \mathrm{H}_{14}$
(c) Gallium metal at $40^{\circ} \mathrm{C}$

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Problem 13

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 14

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

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Problem 15

How would you convert a sample of liquid to vapor without changing the temperature?

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Problem 16

Define the heat of vaporization of a liquid. How is it related to the heat of condensation of that liquid? Using the idea of intermolecular attractions, explain why the process of vaporization is endothermic.

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Problem 17

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 18

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 heat of vaporization versus the number of hydrogen atoms. Explain any trend that you see in terms of intermolecular forces.

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Problem 19

How much thermal energy transfer is required to vaporize $1.0$ metric ton of ammonia? (1 metric ton $=10^{3} \mathrm{~kg}$.) The $\Delta H_{\text {van }}$ for ammonia is $25.1 \mathrm{~kJ} / \mathrm{mol}$.

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Problem 20

The chlorofluorocarbon $\mathrm{CCl}_{3} \mathrm{~F}$ has an enthalpy of vaporization of $24.8 \mathrm{~kJ} / \mathrm{mol}$. To vaporize $1.00 \mathrm{~kg}$ of the compound, how much thermal energy transfer is required?

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Problem 21

The molar enthalpy of vaporization of methanol is $38.0 \mathrm{~kJ} / \mathrm{mol}$ at $25^{\circ} \mathrm{C}$. How much thermal energy transfer is 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 22

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 $0{ }^{\circ} \mathrm{C}$ ). Assume the enthalpy of vaporization of butane is $24.3 \mathrm{~kJ} / \mathrm{mol}$. If the camp stove fuel tank contains $190 .$ g liquid $\mathrm{C}_{4} \mathrm{H}_{10}$, how much thermal energy transfer is required to vaporize all of the butane?

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Problem 23

Mercury is highly toxic. Although it is a liquid at room temperature, it has a high vapor pressure and a low enthalpy of vaporization ( $294 \mathrm{~J} / \mathrm{g}$ ). What quantity of ther-

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Problem 24

Rationalize 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 25

Briefly explain the variations in the boiling points in this table. In your discussion be sure to mention the types of intermolecular forces involved.

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Problem 26

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

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Problem 27

Methanol, $\mathrm{CH}_{3} \mathrm{OH}$, has a normal boiling point of $64.7^{\circ} \mathrm{C}$ and a vapor pressure of $100 \mathrm{~mm} \mathrm{Hg}$ 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 $100 \mathrm{~mm} \mathrm{Hg}$ 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 28

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 29

The lowest sea-level barometric pressure ever recorded was $25.90$ 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? Remember that 1 atm is $760 \mathrm{~mm}$ (29.92 in) $\mathrm{Hg}$, and use Figure $11.5$.

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Problem 30

The highest mountain in the western hemisphere is Mt. Aconcagua, in the central Andes of Argentina $(22,834 \mathrm{ft})$. If atmospheric pressure decreases at a rate of $3.5$ millibar every $100 \mathrm{ft}$, estimate the atmospheric pressure at the top of Mt. Aconcagua, and then estimate from Figure $11.5$ the temperature at which water would boil at the top of the mountain.

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Problem 31

A liquid has a $\Delta H_{\text {vap }}$ 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 32

A liquid has a $\Delta H_{\text {vap }}$ of $44.0 \mathrm{~kJ} / \mathrm{mol}$ and a vapor pressure of $370 \mathrm{~mm} \mathrm{Hg}$ at $90{ }^{\circ} \mathrm{C}$. Calculate the vapor pressure of the liquid at $130^{\circ} \mathrm{C}$.

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Problem 33

The vapor pressure of ethanol, $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}$, at $50.0{ }^{\circ} \mathrm{C}$ is $233 \mathrm{~mm} \mathrm{Hg}$, and its normal boiling point at $1 \mathrm{~atm}$ is $78.3^{\circ} \mathrm{C} .$ Calculate the $\Delta H_{\mathrm{vap}}$ of ethanol.

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Problem 34

Calculate the $\Delta H_{\text {vap }}$ 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 35

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

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Problem 36

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

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Problem 37

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

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Problem 38

The enthalpy of fusion for $\mathrm{H}_{2} \mathrm{O}$ is about $2.5$ times larger than the enthalpy of fusion 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 39

Benzene is an organic liquid that freezes at $5.5^{\circ} \mathrm{C}$ and forms beautiful, feather-like crystals. How much heat transfer occurs when $15.5 \mathrm{~g}$ benzene freezes at $5.5^{\circ} \mathrm{C} ?$ The enthalpy of fusion of benzene is $127 \mathrm{~J} / \mathrm{g}$. If the $15.5 \mathrm{~g}$ sample is remelted, again at $5.5^{\circ} \mathrm{C}$, what quantity of heat transfer is required to c?nvert it to a liquid?

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Problem 40

What is the total quantity of heat energy transfer required to change $0.50 \mathrm{~mol}$ ice at $-5^{\circ} \mathrm{C}$ to $0.50 \mathrm{~mol}$ steam at $100^{\circ} \mathrm{C} ?$

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Problem 41

How much thermal energy is needed to melt a $36.00-\mathrm{g}$ ice cube that is initially at $-10^{\circ} \mathrm{C}$ and bring it to room temperature $\left(20^{\circ} \mathrm{C}\right)$ ? 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 enthalpy of fusion for solid ice is $6.02 \mathrm{~kJ} / \mathrm{mol}$ and the enthalpy of vaporization of liquid water is $40.7 \mathrm{~kJ} / \mathrm{mol}$.

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Problem 42

The chlorofluorocarbon $\mathrm{CCl}_{2} \mathrm{~F}_{2}$ was once used as a refrigerant. What mass of this substance must evaporate to freeze 2 mol water initially at $20^{\circ} \mathrm{C} ?$ The enthalpy of vaporization for $\mathrm{CCl}_{2} \mathrm{~F}_{2}$ is $289 \mathrm{~J} / \mathrm{g} .$ The enthalpy of fusion 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 43

The ions of NaF and $\mathrm{MgO}$ all have the same number of electrons, and the internuclear distances are about the same (235 pm and 212 pm). Why, then, are the melting points of NaF and $\mathrm{MgO}$ so different $\left(992^{\circ} \mathrm{C}\right.$ and $2642^{\circ} \mathrm{C}$, respectively)?

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Problem 44

For the pair of compounds LiF and CsI, tell which compound is expected to have the higher melting point, and briefly explain why.

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Problem 45

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

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Problem 46

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

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Problem 47

Why is solid $\mathrm{CO}_{2}$ called Dry Ice?

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Problem 48

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 49

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 50

From memory, sketch the phase diagram of water. Label all the regions as to the physical state of water. Draw either horizontal (constant pressure) or vertical (constant temperature) paths (i.e., lines with arrows indicating a direction) for these changes of state:
(a) Sublimation
(b) Condensation to a liquid
(c) Melting
(d) Vaporization
(e) Crystallization

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Problem 51

Consult the phase diagram of $\mathrm{CO}_{2}$ in Figure $11.17$. 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 52

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 53

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 54

The molecular structure of water makes it a good solvent for many substances. (a) For what types of substances is water a good solvent? Explain why. (b) For what types of substances is water a poor solvent? Explain why.

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Problem 55

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 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

The presence of large amounts of water (as a lake or ocean) has a moderating effect on the climate near the water. What molecular property of water accounts for this observation?

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Problem 58

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

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Problem 59

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

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Problem 60

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 61

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 62

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 63

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 64

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

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Problem 65

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

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Problem 66

Explain how the volume of a primitive cubic unit cell is related to the radius of the atoms in the cell.

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Problem 67

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

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Problem 68

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

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Problem 69

Using the NaCl structure shown in Figure $11.23$, 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 70

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 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 11.22.)

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Problem 71

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 $3.85 \times 10^{-8} \mathrm{~cm}$, what is the unit cell geometry?

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Problem 72

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

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Problem 73

A primitive cubic unit cell is formed so that the spherical atoms or ions just touch one another along the edge. Prove mathematically that the percentage of empty space within the unit cell is $47.6 \%$. (The volume of a sphere is $\frac{4}{3} \pi r^{3}$, where $r$ is the radius of the sphere.)

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Problem 74

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. The cubic unit cell is $600 \mathrm{pm}$ 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.

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Problem 75

Explain why diamond is denser than graphite.

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Problem 76

Determine, by looking up data in a reference such as the Handbook of Chemistry and Pbysics, 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 77

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

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Problem 78

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 its 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 79

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 unitcell size of the 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 80

The first-order Bragg reflection $(n=1)$ from an aluminum crystal for X-rays with a wavelength of $154 \mathrm{pm}$ is $19.3^{\circ}$. What is the spacing between the planes of aluminum atoms?

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Problem 81

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

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Problem 82

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

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Problem 83

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

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Problem 84

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 85

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

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Problem 86

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

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Problem 87

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

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Problem 88

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

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Problem 89

What is the main technological or economic barrier to the widespread use of superconductors?

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Problem 90

What are the two main chemical reactions involved in the
production of electronic-grade silicon? Identify the elements being reduced and being oxidized.

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Problem 91

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 92

What is the process of doping, as applied to semiconductors? Why are Group $3 \mathrm{~A}$ and Group 5 a elements used to dope silicon?

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Problem 93

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

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Problem 94

Define the term "amorphous."

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Problem 95

What makes a glass different from a solid such as NaCl? Under what conditions could NaCl become glass-like?

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Problem 96

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 97

Give two examples of (a) oxide ceramics and (b) nonox-
ide ceramics.

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Problem 98

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 enthalpy of vaporization 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. How
much thermal energy is evolved when $10.0 \mathrm{~g} \mathrm{CCl}_{2} \mathrm{~F}_{2}$ is cooled from $40^{\circ} \mathrm{C}$ to $-40^{\circ} \mathrm{C} ?$

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Numerade Educator

Problem 99

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 enthalpy of vaporization 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 $4.7 \mathrm{Jg}^{-1} \mathrm{~K}^{-1}$, respectively. How much thermal energy transfer is required to raise the temperature of $10.0 \mathrm{~kg}$ liquid ammonia from $-50.0^{\circ} \mathrm{C}$ to $-33.4^{\circ} \mathrm{C}$, and then to $0.0^{\circ} \mathrm{C} ?$

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 100

Potassium chloride and rubidium chloride both have the sodium chloride structure. X-ray diffraction experiments indicate that their cubic unit cell dimensions are $629 \mathrm{pm}$ and $658 \mathrm{pm}$, respectively.
(i) One mol 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 mol KCl and 1 mol RbCl are heated until the entire mixture is molten and then cooled to room

Shazia Naz

Numerade Educator

Problem 101

Sulfur dioxide, $\mathrm{SO}_{2}$, is found in polluted air.
(a) What types of forces are responsible for binding $\mathrm{SO}_{2}$ molecules to one another in the solid or liquid phase?
(b) Using the information below, place the compounds listed in order of increasing intermolecular attractions.

Aadit Sharma

Numerade Educator

Problem 102

Refer to Figure $11.5$ when answering these questions.
(a) What is the equilibrium vapor pressure for ethyl alcohol at room temperature?
(b) At what temperature $\left({ }^{\circ} \mathrm{C}\right)$ does diethyl ether have an equilibrium vapor pressure of $400 \mathrm{~mm} \mathrm{Hg}$ ?
(c) If water is boiling at a temperature of $95^{\circ} \mathrm{C}$, what is the atmospheric pressure?
(d) At $200 \mathrm{~mm} \mathrm{Hg}$ and $60{ }^{\circ} \mathrm{C}$, which of the three substances are gases?
(e) If you put a couple of drops of each substance on your hand, which would immediately evaporate, and which would remain as a liquid?
(f) Which of the three substances has the greatest intermolecular attractions?

Aadit Sharma

Numerade Educator

Problem 103

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 104

Butane is a gas at room temperature; however, if you look closely at a butane lighter you see it contains liquid butane. How is this possible?

Aadit Sharma

Numerade Educator

Problem 105

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 106

Examine the nanoscale diagrams and the phase diagram below. Match each particulate diagram ( 1 through 7 ) to its corresponding point (A through $\mathrm{H}$ ) on the phase diagram.

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 107

Consider the phase diagram below. 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 108

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.)

Crystal Wang

Numerade Educator

Problem 109

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 110

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 111

We hear reports from weather forecasters of "relative humidity," the ratio of the partial pressure of water in the air to the equilibrium vapor pressure of water at the same temperature. (The vapor pressure of water at $32.2^{\circ} \mathrm{C}$ is $36 \mathrm{~mm} \mathrm{Hg}$.)
(a) On a sticky, humid day, the relative humidity may reach $90 \%$ with a temperature of $90 .{ }^{\circ} \mathrm{F}\left(32.2{ }^{\circ} \mathrm{C}\right)$. What is the partial pressure of water under these conditions? How many moles per liter are present in the air? How many water molecules per $\mathrm{cm}^{3}$ are present in such air?
(b) On a day in a desert, the relative humidity may be $5 . \%$ with the same temperature of $90 .{ }^{\circ} \mathrm{F}$. How many water molecules per $\mathrm{cm}^{3}$ are present in such air?

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 112

(a) In the diagram below, for a substance going from point $\mathrm{F}$ (initial state) to point $\mathrm{G}$ (final state), what changes in phase occur?
(b) What does point A represent?
(c) What does the curve from point $\mathrm{A}$ to point $\mathrm{B}$ represent?

Aadit Sharma

Numerade Educator

Problem 113

Suppose that liquid A has stronger intermolecular forces than liquid $\mathrm{B}$ at room temperature.
(a) Which substance will have the greater surface tension?
(b) Which substance will have the greater vapor pressure?
(c) Which substance will have the greater viscosity?

Aadit Sharma

Numerade Educator

Problem 114

Use the vapor pressure curves shown in the figure below 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 $600 \mathrm{~mm} \mathrm{Hg}$ ?
(d) What are the normal boiling points of these three liquids?

Aadit Sharma

Numerade Educator

Problem 115

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.

Pronoy Sinha

Numerade Educator

Problem 116

Calculate the boiling point of water at $24 \mathrm{~mm} \mathrm{Hg}$. (The $\Delta H_{\text {van }}$ of water is $40.7 \mathrm{~kJ} / \mathrm{mol}$.

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 117

What is the concentration in mol/L of water vapor in air at
$25^{\circ} \mathrm{C}$ at saturation?

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 118

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.

Pronoy Sinha

Numerade Educator

Problem 119

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

Aadit Sharma

Numerade Educator

Problem 120

Copper is an important metal in the U.S. economy. Most of it is mined in the form of the mineral chalcopyrite, $\mathrm{CuFeS}_{2}$
(a) To obtain one metric ton (1000. kilograms) of copper metal, how many metric tons of chalcopyrite would you have to mine?
(b) If the sulfur in chalcopyrite is converted to $\mathrm{SO}_{2}$, how many metric tons of the gas would you get from one metric ton of chalcopyrite?
(c) Copper crystallizes as a face-centered cubic lattice. Knowing that the density of copper is $8.95 \mathrm{~g} / \mathrm{cm}^{3}$, calculate the radius of the copper atom.

Farhana Sharmin

Farhana Sharmin

Numerade Educator

Problem 121

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:
$\mathrm{C}_{6} \mathrm{H}_{6}\left(80{ }^{\circ} \mathrm{C}\right), \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{Cl}\left(131^{\circ} \mathrm{C}\right)$
$\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{Br}\left(156{ }^{\circ} \mathrm{C}\right), \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}\left(182{ }^{\circ} \mathrm{C}\right)$

Aadit Sharma

Numerade Educator