A solid is soft and has a low melting point (below 100°C).

When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 408.7 pm. Calculate the density of silver.

Shown here is a zinc oxide unit cell. What is the formula of the compound?

Give some evidence that all atoms and molecules exert attractive forces on one another.

The vapor pressure of ethanol is 100 mmHg at \(34.9^{\circ} \mathrm{C}\). What is its vapor pressure at \(63.5^{\circ} \mathrm{C}\)? (\(\Delta H_{\mathrm {vap }}\) for ethanol is 39.3 kJ/mol.)

If you lived in Alaska, which of the following natural gases would you keep in an outdoor storage tank in winter? Explain why. methane \((\mathrm{CH}_{4})\), propane \((\mathrm{C}_{3} \mathrm{H}_{8})\), or butane \((\mathrm{C}_{4} \mathrm{H}_{10})\)

Calculate the heat released when 68.0 g of steam at \(124^{\circ} \mathrm{C}\) is converted to water at \(45^{\circ} \mathrm{C}\)

What kind of attractive forces must be overcome in order to (a) melt ice, (b) boil molecular bromine, (c) melt solid iodine, and (d) dissociate \(\mathrm{F}_{2}\) into \(\mathrm{F}\) atoms?

The following compounds have the same molecular formulas \(\left(\mathrm{C}_{4} \mathrm{H}_{10}\right)\). Which one would you expect to have a higher boiling point?

Outdoor water pipes have to be drained or insulated in winter in a cold climate. Why?

Predict which of the following liquids has greater surface tension: ethanol \(\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\right)\) or dimethyl ether \(\left(\mathrm{CH}_{3} \mathrm{OCH}_{3}\right)\).

Predict the viscosity of ethylene glycol relative to that of ethanol and glycerol (see Table 11.3). \(\mathrm{CH}_{2}?\mathrm{OH}\) ? \(\mathrm{CH}_{2}?\mathrm{OH}\) ethylene glycol

Europium crystallizes in a body-centered cubic lattice (the \(\mathrm{Eu}\) atoms occupy only the lattice points). The density of \(\mathrm{Eu}\) is 5.26 \(5.26\mathrm{\ g}/\mathrm{cm}^3\). Calculate the unit cell edge length in pm.

Crystalline silicon has a cubic structure. The unit cell edge length is 543 pm. The density of the solid is \(2.33\mathrm{\ g}/\mathrm{cm}^3\). Calculate the number of \(\mathrm{Si}\) atoms in one unit cell.

A face-centered cubic cell contains \(8 \ \mathrm{X}\) atoms at the corners of the cell and \(6 \ \mathrm{Y}\) atoms at the faces. What is the empirical formula of the solid?

Which of the following are molecular solids and which are covalent solids? \(\mathrm{Se}_{8}\), \(\mathrm{HBr}\), \(\mathrm{Si}\), \(\mathrm{CO}_{2}\), \(\mathrm{C}\), \(\mathrm{P}_{4} \mathrm{O}_{6}\), \(\mathrm{SiH}_{4}\)

Classify the solid state of the following substances as ionic crystals, covalent crystals, molecular crystals, or metallic crystals: (a) \(\mathrm{CO}_{2}\), (b) \(\mathrm{B}_{12}\), (c) \(\mathrm{S}_{8}\), (d) \(\mathrm{KBr}\), (e) \(\mathrm{Mg}\), (f) \(\mathrm{SiO}_{2}\), (g) \(\mathrm{LiCl}\), (h) \(\mathrm{Cr}\).

Explain why diamond is harder than graphite. Why is graphite an electrical conductor but diamond is not?

Define boiling point. How does the boiling point of a liquid depend on external pressure? Referring to Table 5.3, what is the boiling point of water when the external pressure is 187.5 mmHg?

As a liquid is heated at constant pressure, its temperature rises. This trend continues until the boiling point of the liquid is reached. No further rise in temperature of the liquid can be induced by heating. Explain.

What is critical temperature? What is the significance of critical temperature in liquefaction of gases?

The molar heats of fusion and sublimation of molecular iodine are 15.27 kJ/mol and 62.30 kJ/mol, respectively. Estimate the molar heat of vaporization of liquid iodine.

Freeze-dried coffee is prepared by freezing brewed coffee and then removing the ice component with a vacuum pump. Describe the phase changes taking place during these processes.

A student hangs wet clothes outdoors on a winter day when the temperature is \(-15^{\circ} \mathrm{C}\). After a few hours, the clothes are found to be fairly dry. Describe the phase changes in this drying process.

The boiling point and freezing point of sulfur dioxide are \(-10^{\circ} \mathrm{C}\) and \(-72.7^{\circ} \mathrm{C}\) (at 1 atm), respectively. The triple point is \(-75.5^{\circ} \mathrm{C}\) and \(1.65\times10^{-3}\mathrm{\ atm}\), and its critical point is at \(157^{\circ} \mathrm{C}\) and 78 atm. On the basis of this information, draw a rough sketch of the phase diagram of \(\mathrm{SO}_{2}\).

A phase diagram of water is shown at the end of this problem. Label the regions. Predict what would happen as a result of the following changes: (a) Starting at \(\mathrm{A}\), we raise the temperature at constant pressure. (b) Starting at \(\mathrm{C}\), we lower the temperature at constant pressure. (c) Starting at \(\mathrm{B}\), we lower the pressure at constant temperature.

The fluorides of the second-period elements and their melting points are: LiF, 845°C; \(B e F_{2}\), 800°C; \(B F_{3}\), ?126.7°C; \(C F_{4}\), ?184°C; \(N F_{3}\), ?206.6°C; \(O F_{2}\), ?223.8°C; \(F_{2}\), ?219.6°C. Classify the type(s) of intermolecular forces present in each compound.

In 2009, thousands of babies in China became ill from drinking contaminated milk. To falsely boost the milk's protein content, melamine (\(\mathrm{C}_{3} \mathrm{H}_{6} \mathrm{N}_{6}\)) was added to diluted milk because of its high nitrogen composition. Unfortunately, melamine forms a precipitate by hydrogen bonding with cyanuric acid (\(\mathrm{C}_{3} \mathrm{H}_{3} \mathrm{N}_{3} \mathrm{O}_{3}\)), another contaminant present. The resulting stonelike particles caused severe kidney damage in many babies. Draw the hydrogen-bonded complex formed from these two molecules.

The standard enthalpy of formation of gaseous molecular iodine is 62.4 kJ/mol. Use this information to calculate the molar heat of sublimation of molecular iodine at 25°C.

The following graph shows approximate plots of in P versus 1/T for three compounds: methanol (\(\mathrm{CH}_{3} \mathrm{OH}\)) methyl chloride (\(\mathrm{CH}_{3} \mathrm{Cl}\)), and propane (\(\mathrm{C}_{3} \mathrm{H}_{8}\)), where P is the vapor pressure. Match the lines with these compounds.

Ozone (\(O_{3}\)) is a strong oxidizing agent that can oxidize all the common metals except gold and platinum. A convenient test for ozone is based on its action on mercury. When exposed to ozone, mercury becomes dull looking and sticks to glass tubing (instead of flowing freely through it). Write a balanced equation for the reaction. What property of mercury is altered by its interaction with ozone?

A sample of limestone (\(\mathrm{CaCO}_{3}\)) is heated in a closed vessel until it is partially decomposed. Write an equation for the reaction and state how many phases are present.

Silicon used in computer chips must have an impurity level below \(10^{-9}\) (that is, fewer than one impurity atom for every \(10^{9}\) Si atoms). Silicon is prepared by the reduction of quartz (\(\mathrm{SiO}_{2}\)) with coke (a form of carbon made by the destructive distillation of coal) at about 2000°C: \(\mathrm{SiO}_{2}(s)+2 \mathrm{C}(s) \rightarrow \mathrm{Si}(l)+2 \mathrm{CO}(g)\) Next, solid silicon is separated from other solid impurities by treatment with hydrogen chloride at 350°C to form gaseous trichlorosilane (\(\mathrm{SiCl}_{3} \mathrm{H}\)): \(\mathrm{Si}(\mathrm{s})+3 \mathrm{HCl}(\mathrm{g}) \rightarrow \mathrm{SiCl}_{3} \mathrm{H}(\mathrm{g})+\mathrm{H}_{2}(\mathrm{g})\) Finally, ultrapure Si can be obtained by reversing the above reaction at 1000°C: \(\mathrm{SiCl}_{3} \mathrm{H}(\mathrm{g})+\mathrm{H}_{2}(\mathrm{g}) \rightarrow \mathrm{Si}(\mathrm{s})+3 \mathrm{HCl}(\mathrm{g})\) (a) Trichlorosilane has a vapor pressure of 0.258 atm at ?2°C. What is its normal boiling point? Is trichlorosilane's boiling point consistent with the type of intermolecular forces that exist among its molecules? (The molar heat of vaporization of trichlorosilane is 28.8 kJ/mol.) (b) What types of crystals do Si and \(\mathrm{SiO}_{2}\) form? (c) Silicon has a diamond crystal structure (see Figure 11.28). Each cubic unit cell (edge length a = 543 pm) contains eight Si atoms. If there are \(1.0 \times 10^{13}\) boron atoms per cubic centimeter in a sample of pure silicon, how many Si atoms are there for every B atom in the sample? Does this sample satisfy the \(10^{-9}\) purity requirement for the electronic grade silicon?

Swimming coaches sometimes suggest that a drop of alcohol (ethanol) placed in an ear plugged with water “draws out the water.” Explain this action from a molecular point of view.

Use the concept of intermolecular forces to explain why the far end of a walking cane rises when one raises the handle.

Why do citrus growers spray their trees with water to protect them from freezing?

The electrical conductance of copper metal decreases with temperature, but that of a \(\mathrm{CuSO}_{4}\) solution increases with temperature. Explain.

Assuming ideal behavior, calculate the density of gaseous HF at its normal boiling point (19.5°C). The experimentally measured density under the same conditions is 3.10 g/L. Account for the discrepancy between your calculated value and the experimental result.

Give an example for each type of intermolecular force. (a) dipole-dipole interaction, (b) dipole-induced dipole interaction, (c) ion-dipole interaction, (d) dispersion forces, (e) van der Waals forces

Name the type(s) of intermolecular forces that exists between molecules (or basic units) in each of the following species: (a) \(\mathrm{LiF}\), (b) \(\mathrm{CH}_{4}\), (c) \(\mathrm{SO}_{2}\)

Which of the following compounds is most likely to exist as a liquid at room temperature: ethane \((\mathrm{C}_{2} \mathrm{H}_{6})\), hydrazine \((\mathrm{N}_{2} \mathrm{H}_{4})\), fluoromethane \((\mathrm{CH}_{3} \mathrm{F})\)?

Explain the term "polarizability." What kind of molecules tend to have high polarizabilities? What is the relationship between polarizability and intermolecular forces?

Which of the following species are capable of hydrogen bonding among themselves? (a) \(\mathrm{H}_{2} \mathrm{S}\), (b) \(\mathrm{C}_{6} \mathrm{H}_{6}\), (c) \(\mathrm{CH}_{3} \mathrm{OH}\)

Why are motorists advised to use more viscous oils for their engines in the summer and less viscous oils in the winter?

X rays of wavelength 0.154 nm are diffracted from a crystal at an angle of \(14.17^{\circ}\). Assuming that n = 1, calculate the distance (in pm) between layers in the crystal.

Why can the X-ray diffraction technique not be used to study molecular structure in a liquid?

How many atoms are in a body-centered cube, assuming that all atoms occupy lattice points?

What physical properties should you consider in comparing the strength of intermolecular forces in solids and in liquids?

A student studies the In \(P\) versus 1/T plots for two organic liquids: methanol \((\mathrm{CH}_{3} \mathrm{OH})\) and dimethyl ether \((\mathrm{CH}_{3} \mathrm{OCH}_{3})\), such as those shown in Figure 11.36. The slopes are \(-2.32\times10^3\mathrm{\ K}\) and \(-4.50\times10^3\mathrm{\ K}\), respectively. How should she assign the \(\Delta H_{\text {vap }}\) values to these two compounds?

Which elements can take part in hydrogen bonding? Why is hydrogen unique in this kind of interaction?

Copper crystallizes in a face-centered cubic lattice (the \(\mathrm{Cu}\) atoms are at the lattice points only). If the density of the metal is \(8.96\mathrm{\ g}/\mathrm{cm}^3\), what is the unit cell edge length in pm?

Which phase diagram corresponds to a substance that will sublime rather than melt as it is heated at 1 atm?

The compounds \(\mathrm{Br}_{2}\) and \(\mathrm{ICl}\) have the same number of electrons, yet \(\mathrm{Br}_{2}\) melts at \(-7.2^{\circ} \mathrm{C}\) and \(\mathrm{ICl}\) melts at \(27.2^{\circ} \mathrm{C}\). Explain.

The binary hydrogen compounds of the Group 4A elements and their boiling points are: \(\mathrm{CH}_{4}\), \(-162^{\circ} \mathrm{C}\); \(\mathrm{SiH}_{4}\), \(-112^{\circ} \mathrm{C}\); \(\mathrm{GeH}_{4}\), \(-88^{\circ} \mathrm{C}\); and \(\mathrm{SnH}_{4}\), \(-52^{\circ} \mathrm{C}\). Explain the increase in boiling points from \(\mathrm{CH}_{4}\) to \(\mathrm{SnH}_{4}\).

List the types of intermolecular forces that exist between molecules (or basic units) in each of the following species: (a) benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)\), (b) \(\mathrm{CH}_{3} \mathrm{Cl}\), (c) \(\mathrm{PF}_{3}\), (d) \(\mathrm{NaCl}\), (e) \(\mathrm{CS}_{2}\).

Ammonia is both a donor and an acceptor of hydrogen in hydrogen-bond formation. Draw a diagram showing the hydrogen bonding of an ammonia molecule with two other ammonia molecules.

Which of the following species are capable of hydrogen-bonding among themselves? (a) \(\mathrm{C}_{2} \mathrm{H}_{6}\), (b) \(\mathrm{HI}\), (c) \(\mathrm{KF}\), (d) \(\mathrm{BeH}_{2}\), (e) \(\mathrm{CH}_{3} \mathrm{COOH}\)

Arrange the following in order of increasing boiling point: \(\mathrm{RbF}\), \(\mathrm{CO}_{2}\), \(\mathrm{CH}_{3} \mathrm{OH}\), \(\mathrm{CH}_{3} \mathrm{Br}\). Explain your reasoning.

Diethyl ether has a boiling point of \(34.5^{\circ} \mathrm{C}\), and 1-butanol has a boiling point of \(117^{\circ} \mathrm{C}\): Both of these compounds have the same numbers and types of atoms. Explain the difference in their boiling points.

Which member of each of the following pairs of substances would you expect to have a higher boiling point? (a) \(\mathrm{O}_{2}\) and \(\mathrm{Cl}_{2}\), (b) \(\mathrm{SO}_{2}\) and \(\mathrm{CO}_{2}\), (c) \(\mathrm{HF}\) and \(\mathrm{HI}\)

Which substance in each of the following pairs would you expect to have the higher boiling point? Explain why. (a) \(\mathrm{Ne}\) or \(\mathrm{Xe}\), (b) \(\mathrm{CO}_{2}\) or \(\mathrm{CS}_{2}\), (c) \(\mathrm{CH}_{4}\) or \(\mathrm{Cl}_{2}\), (d) \(\mathrm{F}_{2}\) or \(\mathrm{LiF}\), (e) \(\mathrm{NH}_{3}\) or \(\mathrm{PH}_{3}\)

Explain in terms of intermolecular forces why (a) \(\mathrm{NH}_{3}\) has a higher boiling point than \(\mathrm{CH}_{4}\) and (b) \(\mathrm{KCl}\) has a higher melting point than \(\mathrm{I}_{2}\).

Explain why liquids, unlike gases, are virtually incompressible.

What is surface tension? What is the relationship between intermolecular forces and surface tension? How does surface tension change with temperature?

Despite the fact that stainless steel is much denser than water, a stainless-steel razor blade can be made to float on water. Why?

Use water and mercury as examples to explain adhesion and cohesion.

A glass can be filled slightly above the rim with water. Explain why the water does not overflow.

Draw diagrams showing the capillary action of (a) water and (b) mercury in three tubes of different radii.

Why does the viscosity of a liquid decrease with increasing temperature?

What is viscosity? What is the relationship between intermolecular forces and viscosity?

Why is ice less dense than water?

Define the following terms: crystalline solid, lattice point, unit cell, coordination number, closest packing.

Describe the geometries of the following cubic cells: simple cubic, body-centered cubic, face-centered cubic. Which of these structures would give the highest density for the same type of atoms? Which the lowest?

Classify the solid states in terms of crystal types of the elements in the third period of the periodic table. Predict the trends in their melting points and boiling points.

The melting points of the oxides of the third-period elements are given in parentheses: \(\mathrm{Na}_{2} \mathrm{O}\) \(\left(1275^{\circ} \mathrm{C}\right)\), \(\mathrm{MgO}\) \(\left(2800^{\circ} \mathrm{C}\right)\), \(\mathrm{Al}_{2} \mathrm{O}_{3}\) \(\left(2045^{\circ} \mathrm{C}\right)\), \(\mathrm{SiO}_{2}\) \(\left(1610^{\circ} \mathrm{C}\right)\), \(\mathrm{P}_{4} \mathrm{O}_{10}\) \(\left(580^{\circ} \mathrm{C}\right)\), \(\mathrm{SO}_{3}\) \(\left(16.8^{\circ} \mathrm{C}\right)\), \(\mathrm{Cl}_{2} \mathrm{O}_{7}\) \(\left(-91.5^{\circ} \mathrm{C}\right)\). Classify these solids in terms of crystal types.

What is the coordination number of each sphere in (a) a simple cubic cell, (b) a body-centered cubic cell, and (c) a face-centered cubic cell? Assume the spheres are all the same.

Calculate the number of spheres that would be found within a simple cubic, a body-centered cubic, and a face-centered cubic cell. Assume that the spheres are the same.

Metallic iron crystallizes in a cubic lattice. The unit cell edge length is 287 pm. The density of iron is \(7.87\mathrm{\ g}/\mathrm{cm}^3\). How many iron atoms are within a unit cell?

Barium metal crystallizes in a body-centered cubic lattice (the \(\mathrm{Ba}\) atoms are at the lattice points only). The unit cell edge length is 502 pm, and the density of the metal is \(3.50\mathrm{\ g}/\mathrm{cm}^3\). Using this information, calculate Avogadro's number. (Hint: First calculate the volume (in \(\mathrm{cm}^3\)) occupied by 1 mole of \(\mathrm{Ba}\) atoms in the unit cells. Next calculate the volume (in \(\mathrm{cm}^3\)) occupied by one \(\mathrm{Ba}\) atom in the unit cell. Assume that \(68 \%\) of the unit cell is occupied by \(\mathrm{Ba}\) atoms.)

Vanadium crystallizes in a body-centered cubic lattice (the \(\mathrm{V}\) atoms occupy only the lattice points). How many \(\mathrm{V}\) atoms are present in a unit cell?

Define X-ray diffraction. What are the typical wavelengths (in nanometers) of X rays (see Figure 7.4)?

Write the Bragg equation. Define every term and describe how this equation can be used to measure interatomic distances.

When X rays of wavelength 0.090 nm are diffracted by a metallic crystal, the angle of first-order diffraction (n = 1) is measured to be \(15.2^{\circ}\). What is the distance (in pm) between the layers of atoms responsible for the diffraction?

The distance between layers in a \(\mathrm{NaCl}\) crystal is 282 pm. X rays are diffracted from these layers at an angle of \(23.0^{\circ}\). Assuming that n = 1, calculate the wavelength of the X rays in nm.

Describe and give examples of the following types of crystals: (a) ionic crystals, (b) covalent crystals, (c) molecular crystals, (d) metallic crystals.

Why are metals good conductors of heat and electricity? Why does the ability of a metal to conduct electricity decrease with increasing temperature?

A solid is hard, brittle, and electrically nonconducting. Its melt (the liquid form of the substance) and an aqueous solution containing the substance conduct electricity. Classify the solid.

A solid is soft and has a low melting point (below \(100^{\circ} \mathrm{C}\)). The solid, its melt, and an aqueous solution containing the substance are all nonconductors of electricity. Classify the solid.

A solid is very hard and has a high melting point. Neither the solid nor its melt conducts electricity. Classify the solid.

What is an amorphous solid? How does it differ from crystalline solid?

Define glass. What is the chief component of glass? Name three types of glass.

What is a phase change? Name all possible changes that can occur among the vapor, liquid, and solid phases of a substance.

What is the equilibrium vapor pressure of a liquid? How is it measured and how does it change with temperature?

Use any one of the phase changes to explain what is meant by dynamic equilibrium.

Define the following terms: (a) molar heat of vaporization, (b) molar heat of fusion, (c) molar heat of sublimation. What are their units?

How is the molar heat of sublimation related to the molar heats of vaporization and fusion? On what law are these relationships based?

What can we learn about the intermolecular forces in a liquid from the molar heat of vaporization?

The greater the molar heat of vaporization of a liquid, the greater its vapor pressure. True or false?

What is the relationship between intermolecular forces in a liquid and the liquid's boiling point and critical temperature? Why is the critical temperature of water greater than that of most other substances?

How do the boiling points and melting points of water and carbon tetrachloride vary with pressure? Explain any difference in behavior of these two substances.

Why is solid carbon dioxide called dry ice?

Wet clothes dry more quickly on a hot, dry day than on a hot, humid day. Explain.

Which of the following phase transitions gives off more heat? (a) 1 mole of steam to 1 mole of water at \(100^{\circ} \mathrm{C}\), or (b) 1 mole of water to 1 mole of ice at \(0^{\circ} \mathrm{C}\).

A beaker of water is heated to boiling by a Bunsen burner. Would adding another burner raise the boiling point of water? Explain.

Calculate the amount of heat (in kJ) required to convert 74.6 g of water to steam at \(100^{\circ} \mathrm{C}\).

How much heat (in kJ) is needed to convert 866 g of ice at \(-10^{\circ} \mathrm{C}\) to steam at \(126^{\circ} \mathrm{C}\)? (The specific heats of ice and steam are \(2.03\mathrm{\ J}/\mathrm{g}\cdot^{\circ}\mathrm{C}\) and \(1.99\mathrm{\ J}/\mathrm{g}\cdot^{\circ}\mathrm{C}\), respectively.)

How is the rate of evaporation of a liquid affected by (a) temperature, (b) the surface area of a liquid exposed to air, (c) intermolecular forces?

Steam at \(100^{\circ} \mathrm{C}\) causes more serious burns than water at \(100^{\circ} \mathrm{C}\). Why?

Vapor pressure measurements at several different temperatures are shown below for mercury. Determine graphically the molar heat of vaporization for mercury. \(t\left({ }^{\circ} \mathrm{C}\right)\) 200 250 300 320 340 \(P\) (mmHg) 17.3 74.4 246.8 376.3 557.9

The vapor pressure of benzene, \(\mathrm{C}_{6} \mathrm{H}_{6}\), is 40.1 mmHg at \(7.6^{\circ} \mathrm{C}\). What is its vapor pressure at \(60.6^{\circ} \mathrm{C}\)? The molar heat of vaporization of benzene is 31.0 kJ/mol.

The vapor pressure of liquid \(\mathrm{X}\) is lower than that of liquid \(\mathrm{Y}\) at \(20^{\circ} \mathrm{C}\), but higher at \(60^{\circ} \mathrm{C}\). What can you deduce about the relative magnitude of the molar heats of vaporization of \(\mathrm{X}\) and \(\mathrm{Y}\)?

Explain why splashing a small amount of liquid nitrogen (b.p. 77 K) is not as harmful as splashing boiling water on your skin.

What is a phase diagram? What useful information can be obtained from the study of a phase diagram?

Explain how water's phase diagram differs from those of most substances. What property of water causes the difference?

The phase diagram of sulfur is shown here. (a) How many triple points are there? (b) Monoclinic and rhombic are two allotropes of sulfur. Which is more stable under atmospheric conditions? (c) Describe what happens when sulfur at 1 atm is heated from \(80^{\circ} \mathrm{C}\) to \(200^{\circ} \mathrm{C}\).

A length of wire is placed on top of a block of ice. The ends of the wire extend over the edges of the ice, and a heavy weight is attached to each end. It is found that the ice under the wire gradually melts, so that the wire slowly moves through the ice block. At the same time, the water above the wire refreezes. Explain the phase changes that accompany this phenomenon.

Which of the following properties indicates very strong intermolecular forces in a liquid? (a) very low surface tension, (b) very low critical temperature, (c) very low boiling point, (d) very low vapor pressure

At \(-35^{\circ} \mathrm{C}\), liquid \(\mathrm{HI}\) has a higher vapor pressure than liquid \(\mathrm{HF}\). Explain.

Referring to Figure 11.41, determine the stable phase of \(\mathrm{CO}_{2}\) at (a) 4 atm and \(-60^{\circ} \mathrm{C}\) and (b) 0.5 atm and \(-20^{\circ} \mathrm{C}\).

Classify the unit cell of molecular iodine.

A \(\mathrm{CO}_{2}\) fire extinguisher is located on the outside of a building in Massachusetts. During the winter months, one can hear a sloshing sound when the extinguisher is gently shaken. In the summertime there is often no sound when it is shaken. Explain. Assume that the extinguisher has no leaks and that it has not been used.

What is the vapor pressure of mercury at its normal boiling point (357°C)?

A flask of water is connected to a powerful vacuum pump. When the pump is turned on. the water begins to boil. After a few minutes, the same water begins to freeze. Eventually, the ice disappears. Explain what happens at each step.

The vapor pressure of a liquid in a closed container depends on which of the following? (a) The volume above the liquid, (b) the amount of liquid present, (c) temperature, (d) intermolecular forces between the molecules in the liquid.

A student is given four solid samples labeled W, X, Y, and Z. All except Z have a metallic luster. She is told that the solids could be gold, lead sulfide, quartz (\(\mathrm{SiO}_{2}\)), and iodine. The results of her investigations are: (a) W is a good electrical conductor; X, Y, and Z are poor electrical conductors. (b) When the solids are hit with a hammer, W flattens out, X shatters into many pieces, Y is smashed into a powder, and Z is cracked. (c) When the solids are heated with a Bunsen burner, Y melts with some sublimation, but X, W, and Z do not melt. (d) In treatment with 6 M \(\mathrm{HNO}_{3}\), X dissolves; there is no effect on W, Y, or Z. On the basis of these test results, identify the solids.

Which of the following statements are false? (a) Dipole-dipole interactions between molecules are greatest if the molecules possess only temporary dipole moments. (b) All compounds containing hydrogen atoms can participate in hydrogen-bond formation. (c) Dispersion forces exist between all atoms, molecules. and ions. (d) The extent of ion-induced dipole interaction depends only on the charge on the ion.

The diagram below shows a kettle of boiling water on a stove. Identify the phases in regions A and B.

The south pole of Mars is covered with dry ice. which partly sublimes during the summer. The \(\mathrm{CO}_{2}\) vapor recondenses in the winter when the temperature drops to 150 K. Given that the heat of sublimation of \(\mathrm{CO}_{2}\) is 25.9 kJ/mol, calculate the atmospheric pressure on the surface of Mars. [Hint: Use Figure 11.41 to determine the normal sublimation temperature of dry ice and Equation (11.5), which also applies to sublimations.]

The properties of gases, liquids, and solids differ in a number of respects. How would you use the kinetic molecular theory (see Section 5.7) to explain the following observations? (a) Ease of compressibility decreases from gas to liquid to solid. (b) Solids retain a definite shape, but gases and liquids do not. (c) For most substances, the volume of a given amount of material increases as it changes from solid to liquid to gas.

Select the substance in each pair that should have the higher boiling point. In each case identify the principal intermolecular forces involved and account briefly for your choice. (a) \(\mathrm{K}_{2} \mathrm{S}\) or \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{N}\), (b) \(\mathrm{Br}_{2}\) or \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}\)

A small drop of oil in water assumes a spherical shape. Explain. (Hint: Oil is made up of nonpolar molecules, which tend to avoid contact with water.)

Under the same conditions of temperature and density, which of the following gases would you expect to behave less ideally: \(\mathrm{CH}_{4}\), \(\mathrm{SO}_{2}\)? Explain.

Determine the final state and its temperature when 150.0 kJ of heat are added to 50.0 g of water at 20°C. The specific heat of steam is \(1.99 \mathrm{~J} / \mathrm{g} \cdot \mathrm{C}\).

The distance between \(\mathrm{Li}^{+}\) and \(\mathrm{Cl}^{-}\) is 257 pm in solid LiCl and 203 pm in a LiCl unit in the gas phase. Explain the difference in the bond lengths.

Heat of hydration, that is, the heat change that occurs when ions become hydrated in solution, is largely due to ion-dipole interactions. The heats of hydration for the alkali metal ions are \(\mathrm{Li}^{+},-520 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{Na}^{+},-405 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{K}^{+},-321 \mathrm{~kJ} / \mathrm{mol}\). Account for the trend in these values.

If water were a linear molecule, (a) would it still be polar, and (b) would the water molecules still be able to form hydrogen bonds with one another?

Calculate the ?H° for the following processes at 25°C: (a) \(B r_{2}(l) \rightarrow B r_{2}(g)\) and (b) \(B r_{2}(g) \rightarrow 2 B r(g)\). Comment on the relative magnitudes of these ?H° values in terms of the forces involved in each case. {Hint: See Table 9.4, and given that \(\Delta H_{f}^{\circ}\left[B r_{2}(g)\right]\) = 30.7 kJ/mol.}

Gaseous or highly volatile liquid anesthetics are often preferred in surgical procedures because once inhaled, these vapors can quickly enter the bloodstream through the alveoli and then enter the brain. Shown here are several common gaseous anesthetics with their boiling points. Based on intermolecular force considerations, explain the advantages of using these anesthetics. (Hint: The brain barrier is made of membranes that have a nonpolar interior region.)

A beaker of water is placed in a closed container. Predict the effect on the vapor pressure of the water when (a) its temperature is lowered, (b) the volume of the container is doubled, (c) more water is added to the beaker.

The phase diagram of helium is shown here. Helium is the only known substance that has two different liquid phases called helium-I and helium-II. (a) What is the maximum temperature at which helium-II can exist? (b) What is the minimum pressure at which solid helium can exist? (c) What is the normal boiling point of helium-I? (d) Can solid helium sublime? (e) How many triple points are there?

Carbon and silicon belong to Group 4A of the periodic table and have the same valence electron configuration (\(n s^{2} n p^{2}\)). Why does silicon dioxide (\(\mathrm{SiO}_{2}\)) have a much higher melting point than carbon dioxide (\(\mathrm{CO}_{2}\))?

A pressure cooker is a sealed container that allows steam to escape when it exceeds a predetermined pressure. How does this device reduce the time needed for cooking?

What are the advantages of cooking the vegetable broccoli with steam instead of boiling it in water?

Provide an explanation for each of the following phenomena: (a) Solid argon (m.p. -189.2°C; b.p. -185.7°C) can be prepared by immersing a flask containing argon gas in liquid nitrogen (b.p. -195.8°C) until it liquefies and then connecting the flask to a vacuum pump. (b) The melting point of cyclohexane (\(C_{6} H_{12}\)) increases with increasing pressure exerted on the solid cyclohexane. (c) Certain high-altitude clouds contain water droplets at - 10°C. (d) When a piece of dry ice is added to a beaker of water, fog forms above the water.

A quantitative measure of how efficiently spheres pack into unit cells is called packing efficiency, which is the percentage of the cell space occupied by the spheres. Calculate the packing efficiencies of a simple cubic cell, a body-centered cubic cell, and a face-centered cubic cell. (Hint: Refer to Figure 11.22 and use the relationship that the volume of a sphere is \(\frac{4}{3} \pi r^{3}\), where r is the radius of the sphere.)

A chemistry instructor performed the following mystery demonstration. Just before the students arrived in class, she heated some water to boiling in an Erlenmeyer flask. She then removed the flask from the flame and closed the flask with a rubber stopper. After the class commenced, she held the flask in front of the students and announced that she could make the water boil simply by rubbing an ice cube on the outside walls of the flask. To the amazement of everyone, it worked. Give an explanation for this phenomenon.

Given the phase diagram of carbon shown, answer the following questions: (a) How many triple points are there and what are the phases that can coexist at each triple point? (b) Which has a higher density, graphite or diamond? (c) Synthetic diamond can be made from graphite. Using the phase diagram, how would you go about making diamond?

The compound dichlorodifluoromethane (\(\mathrm{CCl}_{2} \mathrm{F}_{2}\)) has a normal boiling point of ?30°C, a critical temperature of 112°C, and a corresponding critical pressure of 40 atm. If the gas is compressed to 18 atm at 20°C. will the gas condense? Your answer should be based on a graphical interpretation.

A student heated a beaker of cold water (on a tripod) with a Bunsen burner. When the gas is ignited, she noticed that there was water condensed on the outside of the beaker. Explain what happened.

Iron crystallizes in a body-centered cubic lattice. The cell length as determined by X-ray diffraction is 286.7 pm. Given that the density of iron is 7.874 g/\(\mathrm{cm}^{3}\), calculate Avogadro’s number.

The boiling point of methanol is 65.0°C and the standard enthalpy of formation of methanol vapor is ?201.2 kJ/mol. Calculate the vapor pressure of methanol (in mmHg) at 25°C. (Hint: See Appendix 3 for other thermodynamic data of methanol.)

An alkali metal in the form of a cube of edge length 0.171 cm is vaporized in a 0.843-L container at 1235 K. The vapor pressure is 19.2 mmHg. Identify the metal by calculating the atomic radius in pico-meters and the density. (Hint: You need to consult, Figures 8.5, 11.22, 11.29, and a chemistry handbook.)

A closed vessel of volume 9.6 L contains 2.0 g of water. Calculate the temperature (in °C) at which only half of the water remains in the liquid phase. (See Table 5.3 for vapor pressures of water at different temperatures.)

Is the vapor pressure of a liquid more sensitive to changes in temperature if \(\Delta H_{\mathrm{vap}}\) is small or large?

A sample of water shows the following behavior as it is heated at a constant rate: If twice the mass of water has the same amount of heat transferred to it, which of the following graphs best describes the temperature variation? Note that the scales for all the graphs are the same.

On a summer day the temperature and (relative) humidity were 95°F and 65 percent, respectively, in Florida. What would be the volume of water in a typical student dormitory room if all of the water vapor were condensed to liquid?

Explain the difference between a temporary dipole moment and the permanent dipole moment.

Explain the difference in the melting points of the following compounds: (Hint: Only one of the two can form intramolecular hydrogen bonds.)

The following compounds, listed with their boiling points, are liquid at \(-10^{\circ} \mathrm{C}\): butane, \(-0.5^{\circ} \mathrm{C}\); ethanol, \(78.3^{\circ} \mathrm{C}\); toluene, \(110.6^{\circ} \mathrm{C}\). At \(-10^{\circ} \mathrm{C}\), which of these liquids would you expect to have the highest vapor pressure? Which the lowest? Explain.

Name the kinds of attractive forces that must be overcome in order to (a) boil liquid ammonia, (b) melt solid phosphorus \(\left(\mathrm{P}_{4}\right)\), (c) dissolve \(\mathrm{CsI}\) in liquid \(\mathrm{HF}\), (d) melt potassium metal.

Based on the following properties of elemental boron, classify it as one of the crystalline solids discussed in Section 11.6: high melting point \(\left(2300^{\circ} \mathrm{C}\right)\), poor conductor of heat and electricity, insoluble in water, very hard substance.

The liquid-vapor boundary line in the phase diagram of any substance always stops abruptly at a certain point. Why?

The interionic distance of several alkali halide crystals are: Nacl NaBr NaI KCl KBr KI 282 pm 299 pm 324 pm 315 pm 330 pm 353 pm Plot lattice energy versus the reciprocal interionic distance. How would you explain the plot in terms of the dependence of lattice energy on distance of separation between ions? What law governs this interaction? (For lattice energies, see Table 9.1.)

Which has a greater density, crystalline \(\mathrm{SiO}_{2}\) or amorphous \(\mathrm{SiO}_{2}\)? Why?

Referring to Figure 11.26, determine the number of each type of ion within the unit cells.

A 1.20-g sample of water is injected into an evacuated 5.00-L flask at 65°C. What percentage of the water will be vapor when the system reaches equilibrium? Assume ideal behavior of water vapor and that the volume of liquid water is negligible. The vapor pressure of water at 65°C is 187.5 mmHg.

Argon crystallizes in the face-centered cubic arrangement at 40 K. Given that the atomic radius of argon is 191 pm, calculate the density of solid argon.

What is the origin of dark spots on the inner glass walls of an old tungsten lightbulb? What is the purpose of filling these lightbulbs with argon gas?

Sketch the cooling curves of water from about 110°C to about -10°C. How would you also show the formation of supercooled liquid below 0°C which then freezes to ice? The pressure is at 1 atm throughout the process. The curves need not be drawn quantitatively.

Both calcium and strontium crystallize in face-centered cubic unit cells. Which metal has a greater density?

Estimate the molar heat of vaporization of a liquid whose vapor pressure doubles when the temperature is raised from 85°C to 95°C.