Chemistry

Educators

Problem 1

In terms of their bulk properties, how do liquids and solids differ? How are they similar?

Problem 2

In terms of the kinetic molecular theory, in what ways are liquids similar to solids? In what ways are liquids different from solids?

Problem 3

In terms of the kinetic molecular theory, in what ways are liquids similar to gases? In what ways are liquids different from gases?

Problem 4

Explain why liquids assume the shape of any container into which they are poured, whereas solids are rigid and retain their shape.

Problem 5

What is the evidence that all neutral atoms and molecules exert attractive forces on each other?

Problem 6

Open the PhET States of Matter Simulation (http://openstaxcollege.org/l/16phetvisual) to answer
the following questions:
(a) Select the Solid, Liquid, Gas tab. Explore by selecting different substances, heating and cooling the systems, and changing the state. What similarities do you notice between the four substances for each phase (solid, liquid, gas)? What differences do you notice?
(b) For each substance, select each of the states and record the given temperatures. How do the given temperatures for each state correlate with the strengths of their intermolecular attractions? Explain.
(c) Select the Interaction Potential tab, and use the default neon atoms. Move the Ne atom on the right and observe how the potential energy changes. Select the Total Force button, and move the Ne atom as before. When is the total force on each atom attractive and large enough to matter? Then select the Component Forces button, and move the Ne atom. When do the attractive (van der Waals) and repulsive (electron overlap) forces balance? How does this relate to the potential energy versus the distance between atoms graph? Explain.

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Problem 7

Define the following and give an example of each:
(a) dispersion force
(b) dipole-dipole attraction
(c) hydrogen bond

Problem 8

The types of intermolecular forces in a substance are identical whether it is a solid, a liquid, or a gas. Why then does a substance change phase from a gas to a liquid or to a solid?

Problem 9

Why do the boiling points of the noble gases increase in the order He < Ne < Ar < Kr < Xe?

Problem 10

Neon and HF have approximately the same molecular masses.
(a) Explain why the boiling points of Neon and HF differ.
(b) Compare the change in the boiling points of Ne, Ar, Kr, and Xe with the change of the boiling points of HF, HCl, HBr, and HI, and explain the difference between the changes with increasing atomic or molecular mass.

Problem 11

Arrange each of the following sets of compounds in order of increasing boiling point temperature:
(a) $\mathrm{HCl}, \mathrm{H}_{2} \mathrm{O}, \mathrm{SiH}_{4}$
(b) $\mathrm{F}_{2}, \mathrm{Cl}_{2}, \mathrm{Br}_{2}$
(c) $\mathrm{CH}_{4}, \mathrm{C}_{2} \mathrm{H}_{6}, \mathrm{C}_{3} \mathrm{H}_{8}$
(d) $\mathrm{O}_{2}, \mathrm{NO}, \mathrm{N}_{2}$

Problem 12

The molecular mass of butanol, $\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH},$ is $74.14 ;$ that of ethylene glycol, $\mathrm{CH}_{2} \mathrm{OH},$ is $62.08,$ yet their boiling points are $117.2^{\circ} \mathrm{C}$ and $174^{\circ} \mathrm{C},$ respectively. Explain the reason for the difference.

Problem 13

On the basis of intermolecular attractions, explain the differences in the boiling points of $n$ -butane $\left(-1^{\circ} \mathrm{C}\right)$ and chloreethane $\left(12^{\circ} \mathrm{C}\right),$ which have similar molar masses.

Problem 14

On the basis of dipole moments and/or hydrogen bonding, explain in a qualitative way the differences in the boiling points of acetone $\left(56.2^{\circ} \mathrm{C}\right)$ and 1 -propanol $\left(97.4^{\circ} \mathrm{C}\right),$ which have similar molar masses.

Problem 15

The melting point of $\mathrm{H}_{2} \mathrm{O}(\mathrm{s})$ is $0^{\circ} \mathrm{C}$ . Would you expect the melting point of $\mathrm{H}_{2} \mathrm{S}(\mathrm{s})$ to be $-85^{\circ} \mathrm{C}, 0^{\circ} \mathrm{C},$ or 185 $^{\circ} \mathrm{C} ?$ Explain your answer.

Problem 16

Silane $\left(\mathrm{SiH}_{4}\right),$ phosphine $\left(\mathrm{PH}_{3}\right),$ and hydrogen sulfide $\left(\mathrm{H}_{2} \mathrm{S}\right)$ melt at $-185^{\circ} \mathrm{C},-133^{\circ} \mathrm{C},$ and $-85^{\circ} \mathrm{C}$ respectively. What does this suggest about the polar character and intermolecular attractions of the three compounds?

Problem 17

Explain why a hydrogen bond between two water molecules is weaker than a hydrogen bond between two hydrogen fluoride molecules.

Problem 18

Under certain conditions, molecules of acetic acid, CH3COOH, form “dimers,” pairs of acetic acid molecules held together by strong intermolecular attractions:
Draw a dimer of acetic acid, showing how two CH3COOH molecules are held together, and stating the type of IMF that is responsible.

Problem 19

Proteins are chains of amino acids that can form in a variety of arrangements, one of which is a helix. What kind of IMF is responsible for holding the protein strand in this shape? On the protein image, show the locations of the IMFs that hold the protein together:

Problem 20

The density of liquid $\mathrm{NH}_{3}$ is $0.64 \mathrm{g} / \mathrm{mL} ;$ the density of gaseous $\mathrm{NH}_{3}$ at $\mathrm{STP}$ is 0.0007 $\mathrm{g} / \mathrm{mL}$ . Explain the difference between the densities of these two phases.

Problem 21

Identify the intermolecular forces present in the following solids:
(a) $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}$
(b) $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{3}$
(c) $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{Cl}$

Problem 22

The test tubes shown here contain equal amounts of the specified motor oils. Identical metal spheres were dropped at the same time into each of the tubes, and a brief moment later, the spheres had fallen to the heights indicated in the illustration. Rank the motor oils in order of increasing viscosity, and explain your reasoning:

Problem 23

Although steel is denser than water, a steel needle or paper clip placed carefully lengthwise on the surface of still water can be made to float. Explain at a molecular level how this is possible:

Problem 24

The surface tension and viscosity values for diethyl ether, acetone, ethanol, and ethylene glycol are shown here.
(a) Explain their differences in viscosity in terms of the size and shape of their molecules and their IMFs.
(b) Explain their differences in surface tension in terms of the size and shape of their molecules and their IMFs:

Problem 25

You may have heard someone use the figure of speech “slower than molasses in winter” to describe a process that occurs slowly. Explain why this is an apt idiom, using concepts of molecular size and shape, molecular interactions, and the effect of changing temperature.

Problem 26

It is often recommended that you let your car engine run idle to warm up before driving, especially on cold winter days. While the benefit of prolonged idling is dubious, it is certainly true that a warm engine is more fuel efficient than a cold one. Explain the reason for this.

Problem 27

The surface tension and viscosity of water at several different temperatures are given in this table.
(a) As temperature increases, what happens to the surface tension of water? Explain why this occurs, in terms of molecular interactions and the effect of changing temperature.
(b) As temperature increases, what happens to the viscosity of water? Explain why this occurs, in terms of molecular interactions and the effect of changing temperature.

Problem 28

At $25^{\circ} \mathrm{C},$ how high will water rise in a glass capillary tube with an inner diameter of 0.63 $\mathrm{mm}$ ? Refer to Example 10.4 for the required information.

Problem 29

Water rises in a glass capillary tube to a height of 17 cm. What is the diameter of the capillary tube?

Problem 30

Heat is added to boiling water. Explain why the temperature of the boiling water does not change. What does change?

Problem 31

Heat is added to ice at $0^{\circ} \mathrm{C}$ . Explain why the temperature of the ice does not change. What does change?

Problem 32

What feature characterizes the dynamic equilibrium between a liquid and its vapor in a closed container?

Problem 33

Identify two common observations indicating some liquids have sufficient vapor pressures to noticeably
evaporate?

Problem 34

Identify two common observations indicating some solids, such as dry ice and mothballs, have vapor pressures sufficient to sublime?

Problem 35

What is the relationship between the intermolecular forces in a liquid and its vapor pressure?

Problem 36

What is the relationship between the intermolecular forces in a solid and its melting temperature?

Problem 37

Why does spilled gasoline evaporate more rapidly on a hot day than on a cold day?

Problem 38

Carbon tetrachloride, $\mathrm{CCl}_{4},$ was once used as a dry cleaning solvent, but is no longer used because it is carcinogenic. At $57.8^{\circ} \mathrm{C},$ the vapor pressure of $\mathrm{CCl}_{4}$ is 54.0 $\mathrm{kPa}$ , and its enthalpy of vaporization is 33.05 $\mathrm{kJ} / \mathrm{mol}$ .
Use this information to estimate the normal boiling point for $\mathrm{CCI}_{4}$ .

Problem 39

When is the boiling point of a liquid equal to its normal boiling point?

Problem 40

How does the boiling of a liquid differ from its evaporation?

Problem 41

Use the information in Figure 10.24 to estimate the boiling point of water in Denver when the atmospheric pressure is 83.3 kPa.

Problem 42

A syringe at a temperature of $20^{\circ} \mathrm{C}$ is filled with liquid ether in such a way that there is no space for any vapor. If the temperature is kept constant and the plunger is withdrawn to create a volume that can be occupied by vapor, what would be the approximate pressure of the vapor produced?

Problem 43

Explain the following observations:
(a) It takes longer to cook an egg in Ft. Davis, Texas (altitude, 5000 feet above sea level) than it does in Boston (at sea level).
(b) Perspiring is a mechanism for cooling the body

Problem 44

The enthalpy of vaporization of water is larger than its enthalpy of fusion. Explain why

Problem 45

Explain why the molar enthalpies of vaporization of the following substances increase in the order $\mathrm{CH}_{4} < \mathrm{C}_{2} \mathrm{H}_{6}$ $< \mathrm{C}_{3} \mathrm{H}_{8},$ even though the type of IMF (dispersion) is the same.

Problem 46

Explain why the enthalpies of vaporization of the following substances increase in the order $\mathrm{CH}_{4} < \mathrm{NH}_{3} <$ $\mathrm{H}_{2} \mathrm{O}$ , even though all three substances have approximately the same molar mass.

Problem 47

The enthalpy of vaporization of $\mathrm{CO}_{2}\left(\mathrm{l} \text { ) is } 9.8 \mathrm{kJ} / \mathrm{mol} \text { . Would you expect the enthalpy of vaporization of } \mathrm{CS}_{2}(\mathrm{l})\right.$ to be $28 \mathrm{kJ} / \mathrm{mol}, 9.8 \mathrm{kJ} / \mathrm{mol},$ or $-8.4 \mathrm{kJ} / \mathrm{mol} ?$ Discuss the plausibility of each of these answers.

Problem 48

The hydrogen fluoride molecule, HF, is more polar than a water molecule, $\mathrm{H}_{2} \mathrm{O}$ (for example, has a greater dipole moment), yet the molar enthalpy of vaporization for liquid hydrogen fluoride is luosser than that for water. Explain.

Problem 49

Ethyl chloride (boiling point, $13^{\circ} \mathrm{C} )$ is used anesthetic. When the liquid is sprayed on the skin, it cools the skin enough to freeze and numb it. Explain the cooling effect of liquid ethyl chloride.

Problem 50

Which contains the compounds listed correctly in order of increasing boiling points?
(a) $\mathrm{N}_{2} < \mathrm{CS}_{2} < \mathrm{H}_{2} \mathrm{O} < \mathrm{KCl}$
(b) $\mathrm{H}_{2} \mathrm{O} < \mathrm{N}_{2} < \mathrm{CS}_{2} < \mathrm{KCl}$
(c) $\mathrm{N}_{2} < \mathrm{KCl} < \mathrm{CS}_{2} < \mathrm{H}_{2} \mathrm{O}$
(d) $\mathrm{CS}_{2} < \mathrm{N}_{2} < \mathrm{KCl} < \mathrm{H}_{2} \mathrm{O}$
(e) $\mathrm{KCl} < \mathrm{H}_{2} \mathrm{O} < \mathrm{CS}_{2} < \mathrm{N}_{2}$

Problem 51

How much heat is required to convert 422 g of liquid $\mathrm{H}_{2} \mathrm{O}$ at $23.5^{\circ} \mathrm{C}$ into steam at $150^{\circ} \mathrm{C} ?$

Problem 52

Evaporation of sweat requires energy and thus take excess heat away from the body. Some of the water that you drink may eventually be converted into sweat and evaporate. If you drink a 20-ounce bottle of water that had been in the refrigerator at $3.8^{\circ} \mathrm{C},$ how much heat is needed to convert all of that water into sweat and then to vapor? (Note: Your body temperature is $36.6^{\circ} \mathrm{C} .$ For the purpose of solving this problem, assume that the thermal properties of sweat are the same as for water.)

Problem 53

Titanium tetrachloride, TiCl_ $_{4},$ has a melting point of $-23.2^{\circ} \mathrm{C}$ and has a $\Delta H_{\text { fusion }}=9.37 \mathrm{kJ} / \mathrm{mol}$ .
(a) How much energy is required to melt 263.1 $\mathrm{g} \mathrm{TiCl}_{4} ?$
(b) For TiCl, which will likely have the larger magnitude: $\Delta H_{\text { fusion }}$ or $\Delta H_{\text { vaporization }} ?$ Explain your reasoning.

Problem 54

From the phase diagram for water (Figure 10.31), determine the state of water at:
(a) $35^{\circ} \mathrm{C}$ and 85 $\mathrm{kPa}$
(b) $-15^{\circ} \mathrm{C}$ and 40 $\mathrm{kPa}$
(c) $-15^{\circ} \mathrm{C}$ and 0.1 $\mathrm{kPa}$
(d) $75^{\circ} \mathrm{C}$ and 3 $\mathrm{kPa}$
(e) $40^{\circ} \mathrm{C}$ and 0.1 $\mathrm{kPa}$
(f) $60^{\circ} \mathrm{C}$ and 50 $\mathrm{kPa}$

Problem 55

What phase changes will take place when water is subjected to varying pressure at a constant temperature of $0.005^{\circ} \mathrm{C} ?$ At $40^{\circ} \mathrm{C} ? \mathrm{At}-40^{\circ} \mathrm{C} ?$

Problem 56

Pressure cookers allow food to cook faster because the higher pressure inside the pressure cooker increases the boiling temperature of water. A particular pressure cooker has a safety valve that is set to vent steam if the pressure exceeds 3.4 atm. What is the approximate maximum temperature that can be reached inside this pressure cooker? Explain your reasoning.

Problem 57

From the phase diagram for carbon dioxide in Figure $10.34,$ determine the state of $\mathrm{CO}_{2}$ at:
(a) $20^{\circ} \mathrm{C}$ and 1000 $\mathrm{kPa}$
(b) $10^{\circ} \mathrm{C}$ and 2000 $\mathrm{kPa}$
(c) $10^{\circ} \mathrm{C}$ and 100 $\mathrm{kPa}$
(d) $-40^{\circ} \mathrm{C}$ and 500 $\mathrm{kPa}$
(e) $-80^{\circ} \mathrm{C}$ and 1500 $\mathrm{kPa}$
(f) $-80^{\circ} \mathrm{C}$ and 10 $\mathrm{kPa}$

Problem 58

Determine the phase changes that carbon dioxide undergoes as the pressure changes if the temperature is held at $-50^{\circ} \mathrm{C}$ ? If the temperature is held at $-40^{\circ} \mathrm{C} ?$ At $20^{\circ} \mathrm{C} ?$ (See the phase diagram in Figure $10.34 . )$

Problem 59

Consider a cylinder containing a mixture of liquid carbon dioxide in equilibrium with gaseous carbon dioxide at an initial pressure of 65 atm and a temperature of $20^{\circ} \mathrm{C}$ . Sketch a plot depicting the change in the cylinder pressure with time as gaseous carbon dioxide is released at constant temperature.

Problem 60

Dry ice, $\mathrm{CO}_{2}(\mathrm{s}),$ does not melt at atmospheric pressure. It sublimes at a temperature of $-78^{\circ} \mathrm{C}$ . What is the lowest pressure at which $\mathrm{CO}_{2}(s)$ will melt to give $\mathrm{CO}_{2}(l) ?$ At approximately what temperature will this occur? (See
Figure 10.34 for the phase diagram.)

Problem 61

If a severe storm results in the loss of electricity, it may be necessary to use a clothesline to dry laundry. In many parts of the country in the dead of winter, the clothes will quickly freeze when they are hung on the line. If it does not snow, will they dry anyway? Explain your answer.

Problem 62

Is it possible to liquefy nitrogen at room temperature (about $25^{\circ} \mathrm{C} )$ . Is it possible to liquefy sulfur dioxide at room temperature? Explain your answers.

Problem 63

Elemental carbon has one gas phase, one liquid phase, and two different solid phases, as shown in the phase diagram:
(a) On the phase diagram, label the gas and liquid regions.
(b) Graphite is the most stable phase of carbon at normal conditions. On the phase diagram, label the graphite phase.
(c) If graphite at normal conditions is heated to 2500 $\mathrm{K}$ while the pressure is increased to $10^{10} \mathrm{Pa}, \mathrm{it}$ is converted into diamond. Label the diamond phase.
(d) Circle each triple point on the phase diagram.
(e) In what phase does carbon exist at 5000 $\mathrm{K}$ and $10^{8} \mathrm{Pa}$ ?
(f) If the temperature of a sample of carbon increases from 3000 $\mathrm{K}$ to 5000 $\mathrm{K}$ at a constant pressure of $10^{6} \mathrm{Pa},$ which phase transition occurs, if any?

Problem 64

What types of liquids typically form amorphous solids?

Problem 65

At very low temperatures oxygen, $\mathrm{O}_{2},$ freezes and forms a crystalline solid. Which best describes these crystals?
(a) ionic
(b) covalent network
(c) metallic
(d) amorphous
(e) molecular crystals

Problem 66

As it cools, olive oil slowly solidifies and forms a solid over a range of temperatures. Which best describes the solid?
(a) ionic
(b) covalent network
(c) metallic
(d) amorphous
(e) molecular crystals

Problem 67

Explain why ice, which is a crystalline solid, has a melting temperature of $0^{\circ} \mathrm{C}$ , whereas butter, which is an amorphous solid, sottens over a range of temperatures.

Problem 68

Identify the type of crystalline solid (metallic, network covalent, ionic, or molecular) formed by each of the following substances:
(a) $\mathrm{SiO}_{2}$
(b) $\mathrm{KCl}$
(c) $\mathrm{Cu}$
(d) $\mathrm{CO}_{2}$
(e) C (diamond)
(f) $\mathrm{BaSO}_{4}$
(g) $\mathrm{NH}_{3}$
(h) $\mathrm{NH}_{4} \mathrm{F}$
(i) $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}$

Problem 69

Identify the type of crystalline solid (metallic, network covalent, ionic, or molecular) formed by each of the following substances:
(a) $\mathrm{CaCl}_{2}$
(b) $\mathrm{SiC}$
(c) $\mathrm{N}_{2}$
(d) Fe
(e) C (graphite)
(f) $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}$
(g) $\mathrm{HCl}$
(h) $\mathrm{NH}_{4} \mathrm{NO}_{3}$
(i) $\mathrm{K}_{3} \mathrm{PO}_{4}$

Problem 70

Classify each substance in the table as either a metallic, ionic, molecular, or covalent network solid:

Problem 71

Classify each substance in the table as either a metallic, ionic, molecular, or covalent network solid:

Problem 72

Identify the following substances as ionic, metallic, covalent network, or molecular solids:
Substance A is malleable, ductile, conducts electricity well, and has a melting point of $1135^{\circ} \mathrm{C}$ . Substance $\mathrm{B}$ is brittle, does not conduct electricity as a solid but does when molten, and has a melting point of $2072^{\circ} \mathrm{C}$ . Substance C is very hard, does not conduct electricity, and has a melting point of $3440^{\circ} \mathrm{C}$ . Substance $\mathrm{D}$ is soft, does not conduct electricity, and has a melting point of $185^{\circ} \mathrm{C} .$

Problem 73

Substance A is shiny, conducts electricity well, and melts at $975^{\circ} \mathrm{C}$ . Substance $\mathrm{A}$ is likely a(n):
(a) ionic solid
(b) metallic solid
(c) molecular solid
(d) covalent network solid

Problem 74

Substance $B$ is hard, does not conduct electricity, and melts at $1200^{\circ} \mathrm{C}$ . Substance $\mathrm{B}$ is likely a(n):
(a) ionic solid
(b) metallic solid
(c) molecular solid
(d) covalent network solid

Problem 75

Describe the crystal structure of iron, which crystallizes with two equivalent metal atoms in a cubic unit cell.

Problem 76

Describe the crystal structure of Pt, which crystallizes with four equivalent metal atoms in a cubic unit cell.

Problem 77

What is the coordination number of a chromium atom in the body-centered cubic structure of chromium?

Problem 78

What is the coordination number of an aluminum atom in the face-centered cubic structure of aluminum?

Problem 79

Cobalt metal crystallizes in a hexagonal closest packed structure. What is the coordination number of a cobalt atom?

Problem 80

Nickel metal crystallizes in a cubic closest packed structure. What is the coordination number of a nickel
atom?

Problem 81

Tungsten crystallizes in a body-centered cubic unit cell with an edge length of 3.165 A.
(a) What is the atomic radius of tungsten in this structure?
(b) Calculate the density of tungsten.

Problem 82

Platinum (atomic radius = 1.38 A) crystallizes in a cubic closely packed structure. Calculate the edge length of the face-centered cubic unit cell and the density of platinum.

Problem 83

Barium crystallizes in a body-centered cubic unit cell with an edge length of 5.025 A
(a) What is the atomic radius of barium in this structure?
(b) Calculate the density of barium.

Problem 84

Aluminum (atomic radius = 1.43 A) crystallizes in a cubic closely packed structure. Calculate the edge length of the face-centered cubic unit cell and the density of aluminum.

Problem 85

The density of aluminum is $2.7 \mathrm{g} / \mathrm{cm}^{3} ;$ that of silicon is 2.3 $\mathrm{g} / \mathrm{cm}^{3} .$ Explain why Si has the lower density even though it has heavier atoms.

Problem 86

The free space in a metal may be found by subtracting the volume of the atoms in a unit cell from the volume of the cell. Calculate the percentage of free space in each of the three cubic lattices if all atoms in each are of equal size and touch their nearest neighbors. Which of these structures represents the most efficient packing? That is, which packs with the least amount of unused space?

Problem 87

Cadmium sulfide, sometimes used as a yellow pigment by artists, crystallizes with cadmium, occupying onehalf of the tetrahedral holes in a closest packed array of sulfide ions. What is the formula of cadmium sulfide? Explain your answer.

Problem 88

A compound of cadmium, tin, and phosphorus is used in the fabrication of some semiconductors. It crystallizes with cadmium occupying one-fourth of the tetrahedral holes and tin occupying one-fourth of the tetrahedral holes in a closest packed array of phosphide ions. What is the formula of the compound? Explain your answer.

Problem 89

What is the formula of the magnetic oxide of cobalt, used in recording tapes, that crystallizes with cobalt atoms occupying one-eighth of the tetrahedral holes and one-half of the octahedral holes in a closely packed array of oxide ions?

Problem 90

A compound containing zinc, aluminum, and sulfur crystallizes with a closest-packed array of sulfide ions. Zinc ions are found in one-eighth of the tetrahedral holes and aluminum ions in one-half of the octahedral holes. What is the empirical formula of the compound?

Problem 91

A compound of thallium and iodine crystallizes in a simple cubic array of iodide ions with thallium ions in all of the cubic holes. What is the formula of this iodide? Explain your answer.

Problem 92

Which of the following elements reacts with sulfur to form a solid in which the sulfur atoms form a closestpacked array with all of the octahedral holes occupied: Li, Na, Be, Ca, or Al?

Problem 93

What is the percent by mass of titanium in rutile, a mineral that contains titanium and oxygen, if structure can be described as a closest packed array of oxide ions with titanium ions in one-half of the octahedral holes? What is the oxidation number of titanium?

Problem 94

Explain why the chemically similar alkali metal chlorides NaCl and CsCl have different structures, whereas the chemically different NaCl and MnS have the same structure.

Problem 95

As minerals were formed from the molten magma, different ions occupied the same cites in the crystals.
Lithium often occurs along with magnesium in minerals despite the difference in the charge on their ions. Suggest an explanation.

Problem 96

Rubidium iodide crystallizes with a cubic unit cell that contains iodide ions at the corners and a rubidium ion in the center. What is the formula of the compound?

Problem 97

One of the various manganese oxides crystallizes with a cubic unit cell that contains manganese ions at the corners and in the center. Oxide ions are located at the center of each edge of the unit cell. What is the formula of the compound?

Problem 98

NaH crystallizes with the same crystal structure as NaCl. The edge length of the cubic unit cell of NaH is 4.880 A.
(a) Calculate the ionic radius of $\mathrm{H}^{-}$ . (The ionic radius of $\mathrm{Li}^{+}$ is 0.0 .95 $\mathrm{A} . )$
(b) Calculate the density of NaH.

Problem 99

Thallium(I) iodide crystallizes with the same structure as $\mathrm{CsCl}$ . The edge length of the unit cell of TII is 4.20 $\mathrm{A}$ . Calculate the ionic radius of $\mathrm{TI}^{+}$ . (The ionic radius of $\mathrm{I}^{-}$ is 2.16 $\mathrm{A} . )$

Problem 100

A cubic unit cell contains manganese ions at the corners and fluoride ions at the center of each edge.
(a) What is the empirical formula of this compound? Explain your answer.
(b) What is the coordination number of the $\mathrm{Mn}^{3+}$ ion?
(c) Calculate the edge length of the unit cell if the radius of a Mn' $^{3+}$ ion is 0.65 $\mathrm{A}$ .
(d) Calculate the density of the compound.

Problem 101

What is the spacing between crystal planes that diffract $\mathrm{X}$ -rays with a wavelength of 1.541 $\mathrm{nm}$ at an angle $\theta$ of $15.55^{\circ}$ (first order reflection)?

Problem 102

A diffractometer using $\mathrm{X}$ -rays with a wavelength of 0.2287 nm produced first-order diffraction peak for a crystal angle $\theta=16.21^{\circ} .$ Determine the spacing between the diffracting planes in this crystal.

Problem 103

A metal with spacing between planes equal to 0.4164 nm diffracts $X$ -rays with a wavelength of 0.2879 $\mathrm{nm}$ What is the diffraction angle for the first order diffraction peak?

Problem 104

Gold crystallizes in a face-centered cubic unit cell. The second-order reflection $(\mathrm{n}=2)$ of $\mathrm{X}$ -rays for the planes that make up the tops and bottoms of the unit cells is at $\theta=22.20^{\circ} .$ The wavelength of the X-rays is 1.54 $\mathrm{A}$ . What is the density of metallic gold?

When an electron in an excited molydenum atom falls from the $\mathrm{L}$ to the $\mathrm{K}$ shell, an $\mathrm{X}$ -ray is emitted. These X-rays are diffracted at an angle of $7.75^{\circ}$ by planes with a separation of 2.64 A. What is the difference in energy between the K shell and the L shell in molybdenum assuming a first-order diffraction?