CHAPTER TEN LIQUIDS AND SOLIDS For Review 1. Intermolecular forces are the relatively weak forces between molecules that hold the molecules together in the solid and liquid phases. Intramolecular forces are the forces within a molecule. These are the covalent bonds in a molecule. Intramolecular forces (covalent bonds) are much stronger than intermolecular forces. Dipole forces are the forces that act between polar molecules. The electrostatic attraction between the positive end of one polar molecule and the negative end of another is the dipole force. Dipole forces are generally weaker than hydrogen bonding. Both of these forces are due to dipole moments in molecules. Hydrogen bonding is given a separate name from dipole forces because hydrogen bonding is a particularly strong dipole force. Any neutral molecule that has a hydrogen covalently bonded to N, O, or F exhibits the relatively strong hydrogen bonding intermolecular forces. London dispersion forces are accidental-induced dipole forces. Like dipole forces, London dispersion forces are electrostatic in nature. Dipole forces are the electrostatic forces between molecules having a permanent dipole. London dispersion forces are the electrostatic forces between molecules having an accidental or induced dipole. All covalent molecules (polar and nonpolar) have London dispersion forces, but only polar molecules (those with permanent dipoles) exhibit dipole forces. As the size of a molecule increases, the strength of the London dispersion forces increases. This is because, as the electron cloud about a molecule gets larger, it is easier for the electrons to be drawn away from the nucleus. The molecule is said to be more polarizable. London dispersion (LD) < dipole-dipole < H bonding < metallic bonding, covalent network, ionic. Yes, there is considerable overlap. Consider some of the examples in Exercise 10.98. Benzene (only LD forces) has a higher boiling point than acetone (dipole-dipole forces). Also, there is even more overlap among the stronger forces (metallic, covalent, and ionic). 2. a. Surface tension: the resistance of a liquid to an increase in its surface area. b. Viscosity: the resistance of a liquid to flow. c. Melting point: the temperature (at constant pressure) where a solid converts entirely to a liquid as long as heat is applied. A more detailed definition is the temperature at which the solid and liquid states have the same vapor pressure under conditions where the total pressure is constant. 331 332 CHAPTER 10 LIQUIDS AND SOLIDS d. Boiling point: the temperature (at constant pressure) where a liquid converts entirely to a gas as long as heat is applied. The detailed definition is the temperature at which the vapor pressure of the liquid is exactly equal to the external pressure. e. Vapor pressure: the pressure of the vapor over a liquid at equilibrium. As the strengths of intermolecular forces increase, surface tension, viscosity, melting point and boiling point increase, while vapor pressure decreases. 3. Solid: rigid; has fixed volume and shape; slightly compressible Liquid: definite volume but no specific shape; assumes shape of the container; slightly compressible Gas: no fixed volume or shape; easily compressible 4. a. Crystalline solid: Regular, repeating structure Amorphous solid: Irregular arrangement of atoms or molecules b. Ionic solid: Molecular solid: c. Molecular solid: Made up of discrete covalently bonded molecules held together in the solid phase by weaker forces (LD, dipole, or hydrogen bonds). Discrete, individual molecules Network solid: No discrete molecules; A network solid is one large molecule. The forces holding the atoms together are the covalent bonds between atoms. d. Metallic solid: Completely delocalized electrons, conductor of electricity (cations in a sea of electrons) Network solid: 5. Made up of ions held together by ionic bonding Localized electrons; Insulator or semiconductor Lattice: a three-dimensional system of points designating the positions of the centers of the components of a solid (atoms, ions, or molecules) Unit cell: the smallest repeating unit of a lattice A simple cubic unit cell has an atom, ion or molecule located at the eight corners of a cube. There is one net atom per simple cubic unit cell. Because the atoms in the cubic unit cell are assumed to touch along the cube edge, cube edge = = 2r where r = radius of the atom. A body-centered cubic unit cell has an atom, ion or molecule at the eight corners of a cube and one atom, ion, or molecule located at the center of the cube. There are two net atoms per body-centered cubic unit cell. Because the atoms in the cubic unit cell are assumed to touch along the body diagonal of the cube, body diagonal = 3 = 4r where = cube edge and r = radius of atom. A face-centered cubic unit cell has an atom, ion, or molecule at the eight CHAPTER 10 LIQUIDS AND SOLIDS 333 corners of a cube and an atom, ion, or molecule located at the six faces of the cube. There are four net atoms per face-centered unit cell. Because the atoms in the cubic unit cell are assumed to touch along the face diagonal of the cube, face diagonal = 2 = 4r. 6. Closest packing: the packing of atoms (uniform, hard spheres) in a manner that most efficiently uses the available space with the least amount of empty space. The two types of closest packing are hexagonal closest packing and cubic closest packing. In both closest packed arrangements, the atoms (spheres) are packed in layers. The difference between the two closest packed arrangements is the ordering of the layers. Hexagonal closest packing has the third layer directly over the first layer forming a repeating layer pattern of abab… In cubic closest packing the layer pattern is abcabc… The unit cell for hexagonal closest packing is a hexagonal prism. See Fig. 10.14 for an illustration of the hexagonal prism unit cell. The unit cell for cubic closest packing is the face-centered cubic unit cell. 7. Conductor: The energy difference between the filled and unfilled molecular orbitals is minimal. We call this energy difference the band gap. Because the band gap is minimal, electrons can easily move into the conduction bands (the unfilled molecular orbitals). Insulator: Large band gap; Electrons do not move from the filled molecular orbitals to the conduction bands since the energy difference is large. Semiconductor: Small band gap; The energy difference between the filled and unfilled molecular orbitals is smaller than in insulators, so some electrons can jump into the conduction bands. The band gap, however, is not as small as with conductors, so semiconductors have intermediate conductivity. a. As the temperature is increased, more electrons in the filled molecular orbitals have sufficient kinetic energy to jump into the conduction bands (the unfilled molecular orbitals). b. A photon of light is absorbed by an electron which then has sufficient energy to jump into the conduction bands. c. An impurity either adds electrons at an energy near that of the conduction bands (n-type) or creates holes (unfilled energy levels) at energies in the previously filled molecular orbitals (p-type). Both n-type and p-type semiconductors increase conductivity by creating an easier path for electrons to jump from filled to unfilled energy levels. In conductors, electrical conductivity is inversely proportional to temperature. Increases in temperature increase the motions of the atoms, which gives rise to increased resistance (decreased conductivity). In a semiconductor, electrical conductivity is directly proportional to temperature. An increase in temperature provides more electrons with enough kinetic energy to jump from the filled molecular orbitals to the conduction bands, increasing conductivity. To produce an n-type semiconductor, dope Ge with a substance that has more than 4 valence electrons, e.g., a group 5A element. Phosphorus or arsenic are two substances which will produce n-type semiconductors when they are doped into germanium. To produce a p-type semiconductor, dope Ge with a substance that has fewer than 4 valence electrons, e.g., a 334 CHAPTER 10 LIQUIDS AND SOLIDS group 3A element. Gallium or indium are two substances which will produce p-type semiconductors when they are doped into germanium. 8. The structures of most binary ionic solids can be explained by the closest packing of spheres. Typically, the larger ions, usually the anions, are packed in one of the closest packing arrangements, and the smaller cations fit into holes among the closest packed anions. There are different types of holes within the closest packed anions which are determined by the number of spheres that form them. Which of the three types of holes are filled usually depends on the relative size of the cation to the anion. Ionic solids will always try to maximize electrostatic attractions among oppositely charged ions and minimize the repulsions among ions with like charges. The structure of sodium chloride can be described in terms of a cubic closest packed array of Cl ions with Na+ ions in all of the octahedral holes. An octahedral hole is formed between 6 Cl anions. The number of octrahedral holes is the same as the number of packed ions. So in the face-centered unit cell of sodium chloride, there are 4 net Cl ions and 4 net octahedral holes. Because the stoichiometry dictates a 1:1 ratio between the number of Cl anions and Na+ cations, all of the octahedral holes must be filled with Na+ ions. In zinc sulfide, the sulfide anions also occupy the lattice points of a cubic closest packing arrangement. But instead of having the cations in octahedral holes, the Zn2+ cations occupy tetrahedral holes. A tetrahedral hole is the empty space created when four spheres are packed together. There are twice as many tetrahedral holes as packed anions in the closest packed structure. Therefore, each face-centered unit cell of sulfide anions contains 4 net S2 ions and 8 net tetrahedral holes. For the 1:1 stoichiometry to work out, only one-half of the tetrahedral holes are filled with Zn2+ ions. This gives 4 S2 ions and 4 Zn2+ ions per unit cell for an empirical formula of ZnS. 9. a. Evaporation: process where liquid molecules escape the liquid’s surface to form a gas. b. Condensation: process where gas molecules hit the surface of a liquid and convert to a liquid. c. Sublimation: process where a solid converts directly to a gas without passing through the liquid state. d. Boiling: the temperature and pressure at which a liquid completely converts to a gas as long as heat is applied. e. Melting: temperature and pressure at which a solid completely converts to a liquid as long as heat is applied. f. Enthalpy of vaporization (Hvap): the enthalpy change that occurs at the boiling point when a liquid converts into a gas. g. Enthalpy of fusion (Hfus): the enthalpy change that occurs at the melting point when a solid converts into a liquid. CHAPTER 10 LIQUIDS AND SOLIDS 335 h. Heating curve: a plot of temperature versus time as heat is applied at a constant rate to some substance. Fusion refers to a solid converting to a liquid, and vaporization refers to a liquid converting to a gas. Only a fraction of the hydrogen bonds in ice are broken in going from the solid phase to the liquid phase. Most of the hydrogen bonds in water are still present in the liquid phase and must be broken during the liquid to gas phase transition. Thus, the enthalpy of vaporization is much larger than the enthalpy of fusion because more intermolecular forces are broken during the vaporization process. A volatile liquid is one that evaporates relatively easily. Volatile liquids have large vapor pressures because the intermolecular forces that prevent evaporation are relatively weak. 10. See Fig. 10.49 and 10.52 for the phase diagrams of H2O and CO2. Most substances exhibit only three different phases: solid, liquid, and gas. This is true for H2O and CO2. Also typical of phase diagrams is the positive slopes for both the liquid-gas equilibrium line and the solidgas equilibrium line. This is also true for both H2O and CO2. The solid-liquid equilibrium line also generally has a positive slope. This is true for CO2, but not for H2O. In the H2O phase diagram, the slope of the solid-liquid line is negative. The determining factor for the slope of the solid-liquid line is the relative densities of the solid and liquid phases. The solid phase is denser than the liquid phase in most substances; for these substances, the slope of the solidliquid equilibrium line is positive. For water, the liquid phase is denser than the solid phase which corresponds to a negative sloping solid-liquid equilibrium line. Another difference between H2O and CO2 is the normal melting points and normal boiling points. The term normal just dictates a pressure of 1 atm. H2O has a normal melting point (0C) and a normal boiling point (100C), but CO2 does not. At 1 atm pressure, CO2 only sublimes (goes from the solid phase directly to the gas phase). There are no temperatures at 1 atm for CO 2 where the solid and liquid phases are in equilibrium or where the liquid and gas phases are in equilibrium. There are other differences, but those discussed above are the major ones. The relationship between melting points and pressure is determined by the slope of the solidliquid equilibrium line. For most substances (CO2 included), the positive slope of the solidliquid line shows a direct relationship between the melting point and pressure. As pressure increases, the melting point increases. Water is just the opposite since the slope of the solidliquid line in water is negative. Here the melting point of water is inversely related to the pressure. For boiling points, the positive slope of the liquid-gas equilibrium line indicates a direct relationship between the boiling point and pressure. This direct relationship is true for all substances including H2O and CO2. The critical temperature for a substance is defined as the temperature above which the vapor cannot be liquefied no matter what pressure is applied. The critical temperature, like the boiling point temperature, is directly related to the strength of the intermolecular forces. Since H2O exhibits relatively strong hydrogen bonding interactions and CO2 only exhibits London dispersion forces, one would expect a higher critical temperature for H2O than for CO2. 336 CHAPTER 10 LIQUIDS AND SOLIDS Questions 12. C25H52 has the stronger intermolecular forces because it has the higher boiling point. Even though C25H52 is nonpolar, it is so large that its London dispersion forces are much stronger than the sum of the London dispersion and hydrogen bonding interactions found in H2O. 13. Atoms have an approximately spherical shape (on the average). It is impossible to pack spheres together without some empty space among the spheres. 14. Critical temperature: Critical pressure: The temperature above which a liquid cannot exist, i.e., the gas cannot be liquified by increased pressure. The pressure that must be applied to a substance at its critical temperature to produce a liquid. The kinetic energy distribution changes as one raises the temperature (T4 > Tc > T3 > T2 > T1). At the critical temperature, Tc, all molecules have kinetic energies greater than the intermolecular forces, F, and a liquid can't form. Note: The distributions above are not to scale. 15. Evaporation takes place when some molecules at the surface of a liquid have enough energy to break the intermolecular forces holding them in the liquid phase. When a liquid evaporates, the molecules that escape have high kinetic energies. The average kinetic energy of the remaining molecules is lower, thus, the temperature of the liquid is lower. 16. A crystalline solid will have the simpler diffraction pattern because a regular, repeating arrangement is necessary to produce planes of atoms that will diffract the X-rays in regular patterns. An amorphous solid does not have a regular repeating arrangement and will produce a complicated diffraction pattern. 17. An alloy is a substance that contains a mixture of elements and has metallic properties. In a substitutional alloy, some of the host metal atoms are replaced by other metal atoms of similar size, e.g., brass, pewter, plumber’s solder. An interstitial alloy is formed when some of the interstices (holes) in the closest packed metal structure are occupied by smaller atoms, e.g., carbon steels. 18. Equilibrium: There is no change in composition; the vapor pressure is constant. Dynamic: Two processes, vapor → liquid and liquid → vapor, are both occurring but with equal rates so the composition of the vapor is constant. CHAPTER 10 19. LIQUIDS AND SOLIDS 337 a. As the strength of the intermolecular forces increase, the rate of evaporation decreases. b. As temperature increases, the rate of evaporation increases. c. As surface area increases, the rate of evaporation increases. 20. C2H5OH(l) → C2H5OH(g) is an endothermic process. Heat is absorbed when liquid ethanol vaporizes; the internal heat from the body provides this heat which results in the cooling of the body. 21. Sublimation will occur allowing water to escape as H2O(g). 22. The phase change, H2O(g) → H2O(l), releases heat that can cause additional damage. Also steam can be at a temperature greater than 100°C. 23. The strength of intermolecular forces determines relative boiling points. The types of intermolecular forces for covalent compounds are London dispersion forces, dipole forces, and hydrogen bonding. Because the three compounds are assumed to have similar molar mass and shape, the strength of the London dispersion forces will be about equal between the three compounds. One of the compounds will be nonpolar so it only has London dispersion forces. The other two compounds will be polar so they have additional dipole forces and will boil at a higher temperature than the nonpolar compound. One of the polar compounds will have an H covalently bonded to either N, O, or F. This gives rise to the strongest type of covalent intermolecular forces, hydrogen bonding. The compound which hydrogen bonds will have the highest boiling point while the polar compound with no hydrogen bonding will boil at a temperature in the middle of the other compounds. 24. a. Both forms of carbon are network solids. In diamond, each carbon atom is surrounded by a tetrahedral arrangement of other carbon atoms to form a huge molecule. Each carbon atom is covalently bonded to four other carbon atoms. The structure of graphite is based on layers of carbon atoms arranged in fused sixmembered rings. Each carbon atom in a particular layer of graphite is surrounded by three other carbons in a trigonal planar arrangement. This requires sp 2 hybridization. Each carbon has an unhybridized p atomic orbital; all of these p orbitals in each sixmembered ring overlap with each other to form a delocalized electron system. b. Silica is a network solid having an empirical formula of SiO2. The silicon atoms are singly bonded to four oxygens. Each silicon atom is at the center of a tetrahedral arrangement of oxygen atoms which are shared with other silicon atoms. The structure of silica is based on a network of SiO4 tetrahedra with shared oxygen atoms, rather than discrete SiO2 molecules. Silicates closely resemble silica. The structure is based on interconnected SiO4 tetrahedra. However, in contrast to silica, where the O/Si ratio is 2:1, silicates have O/Si ratios greater than 2:1 and contain silicon-oxygen anions. To form a neutral solid silicate, metal cations are needed to balance the charge. In other words, silicates are salts containing metal cations and polyatomic silicon-oxygen anions. 338 CHAPTER 10 LIQUIDS AND SOLIDS When silica is heated above its melting point and cooled rapidly, an amorphous (disordered) solid called glass results. Glass more closely resembles a very viscous solution than it does a crystalline solid. To affect the properties of glass, several different additives are thrown into the mixture. Some of these additives are Na2CO3, B2O3, and K2O, with each compound serving a specific purpose relating to the properties of the glass. 25. a. Both CO2 and H2O are molecular solids. Both have an ordered array of the individual molecules, with the molecular units occupying the lattice points. A difference within each solid lattice is the strength of the intermolecular forces. CO2 is nonpolar and only exhibits London dispersion forces. H2O exhibits the relatively strong hydrogen bonding interactions. The differences in strength is evidenced by the solid phase changes that occur at 1 atm. CO2(s) sublimes at a relatively low temperature of 78C. In sublimation, all of the intermolecular forces are broken. However, H2O(s) doesn’t have a phase change until 0C, and in this phase change from ice to water, only a fraction of the intermolecular forces are broken. The higher temperature and the fact that only a portion of the intermolecular forces are broken are attributed to the strength of the intermolecular forces in H2O(s) as compared to CO2(s). Related to the intermolecular forces are the relative densities of the solid and liquid phases for these two compounds. CO2(s) is denser than CO2(l) while H2O(s) is less dense than H2O(l). For CO2(s), the molecules pack together as close as possible, hence solids are usually more dense than the liquid phase. For H2O, each molecule has two lone pairs and two bonded hydrogen atoms. Because of the equal number of lone pairs and OH bonds, each H2O molecule can form two hydrogen bonding interactions to other H2O molecules. To keep this symmetric arrangement (which maximizes the hydrogen bonding interactions), the H2O(s) molecules occupy positions that create empty space in the lattice. This translates into a smaller density for H2O(s) as compared to H2O(l). b. Both NaCl and CsCl are ionic compounds with the anions at the lattice points of the unit cells and the cations occupying the empty spaces created by anions (called holes). In NaCl, the Cl anions occupy the lattice points of a face-centered unit cell with the Na+ cations occupying the octahedral holes. Octahedral holes are the empty spaces created by six Cl ions. CsCl has the Cl ions at the lattice points of a simple cubic unit cell with the Cs+ cations occupying the middle of the cube. 26. Because silicon carbide is made from Group 4A elements and because it is extremely hard, one would expect SiC to form a covalent network structure similar to diamond. 27. The mathematical equation that relates the vapor pressure of a substance to temperature is: ln Pvap = y ΔH vap 1 +C R T m x +b As shown above, this equation is in the form of the straight line equation. If one plots ln P vap vs. 1/T, the slope of the straight line is Hvap/R. Because Hvap is always positive, the slope of the straight line will be negative. CHAPTER 10 28. LIQUIDS AND SOLIDS 339 The typical phase diagram for a substance shows three phases and has a positive sloping solid-liquid equilibrium line (water is atypical). A sketch of the phase diagram for I2 would look like this: s l g P 90 torr 115oC T Statements a and e are true. For statement a, the liquid phase is always more dense than the gaseous phase (gases are mostly empty space). For statement e, because the triple point is at 90 torr, the liquid phase cannot exist at any pressure less than 90 torr, no matter what the temperature. For statements b, c, and d, examine the phase diagram to prove to yourself that they are false. Exercises Intermolecular Forces and Physical Properties 29. 30. Ionic compounds have ionic forces. Covalent compounds all have London Dispersion (LD) forces, while polar covalent compounds have dipole forces and/or hydrogen bonding forces. For H bonding forces, the covalent compound must have either a NH, OH or FH bond in the molecule. a. LD only b. dipole, LD d. ionic e. LD only (CH4 in a nonpolar covalent compound.) f. g. ionic dipole, LD c. H bonding, LD a. ionic b. LD mostly; C‒F bonds are polar, but polymers like teflon are so large the LD forces are the predominant intermolecular forces. 31. c. LD d. dipole, LD f. g. LD dipole, LD e. H bonding, LD a. OCS; OCS is polar and has dipole-dipole forces in addition to London dispersion (LD) forces. All polar molecules have dipole forces. CO2 is nonpolar and only has LD forces. To predict polarity, draw the Lewis structure and deduce whether the individual bond dipoles cancel. 340 CHAPTER 10 LIQUIDS AND SOLIDS b. SeO2; Both SeO2 and SO2 are polar compounds, so they both have dipole forces as well as LD forces. However, SeO2 is a larger molecule, so it would have stronger LD forces. c. H2NCH2CH2NH2; More extensive hydrogen bonding is possible. d. H2CO; H2CO is polar while CH3CH3 is nonpolar. H2CO has dipole forces in addition to LD forces. e. CH3OH; CH3OH can form relatively strong H bonding interactions, unlike H2CO. 32. Ar exists as individual atoms which are held together in the condensed phases by London dispersion forces. The molecule which will have a boiling point closest to Ar will be a nonpolar substance with about the same molar mass as Ar (39.95 g/mol); this same size nonpolar substance will have about equivalent strength of London dispersion forces. Of the choices, only Cl2 (70.90 g/mol) and F2 (38.00 g/mol) are nonpolar. Because F2 has a molar mass closest to that of Ar, one would expect the boiling point of F2 to be close to that of Ar. 33. a. Neopentane is more compact than n-pentane. There is less surface area contact among neopentane molecules. This leads to weaker LD forces and a lower boiling point. b. HF is capable of H bonding; HCl is not. c. LiCl is ionic, and HCl is a molecular solid with only dipole forces and LD forces. Ionic forces are much stronger than the forces for molecular solids. d. n-Hexane is a larger molecule, so it has stronger LD forces. 34. Ethanol, C2H6O, has 2(4) + 6(1) + 6 = 20 valence electrons. H H H C C H H O H Exhibits H-bonding and London dispersion forces. Dimethyl ether, C2H6O, also has 20 valence electrons. It has a Lewis structure of: H H O H C C H H H Exhibits dipole and London dispersion forces, but no hydrogen bonding since it has no H covalently bonded to the O. CHAPTER 10 LIQUIDS AND SOLIDS 341 Propane, C3H6, has 3(4) + 6(1) = 18 valence electrons. H H H H C C C H H H H Propane only has relatively nonpolar bonds so it is nonpolar. Propane exhibits only London dispersion forces. The three compounds have similar molar mass so the strength of the London dispersion forces will be approximately equivalent. Because dimethyl ether has additional dipole forces, it will boil at a higher temperature than propane. The compound with the highest boiling point is ethanol since it exhibits relatively strong hydrogen bonding forces. The correct matching of boiling points is: ethanol, 78.5C; dimethyl ether, 23C; propane, 42.1C 35. Boiling points and freezing points are assumed directly related to the strength of the intermolecular forces, while vapor pressure is inversely related to the strength of the intermolecular forces. a. HBr; HBr is polar, while Kr and Cl2 are nonpolar. HBr has dipole forces unlike Kr and Cl2. b. NaCl; Ionic forces are much stronger than molecular forces. c. I2; All are nonpolar, so the largest molecule (I2) will have the strongest LD forces and the lowest vapor pressure. d. N2; Nonpolar and smallest, so has the weakest intermolecular forces. e. CH4; Smallest, nonpolar molecule so has the weakest LD forces. f. HF; HF can form relatively strong H bonding interactions unlike the others. g. CH3CH2CH2OH; H bonding, unlike the others, so has strongest intermolecular forces. 36. a. CBr4; Largest of these nonpolar molecules so has strongest LD forces. b. F2; Ionic forces in LiF are much stronger than the covalent forces in F2 and HCl. HCl has dipole forces that the nonpolar F2 does not exhibit; so F2 has the weakest intermolecular forces and the lowest freezing point. c. CH3CH2OH; Can form H bonding interactions unlike the others. d. H2O2; HO-OH structure produces stronger H bonding interactions than HF, so has greatest viscosity. e. H2CO; H2CO is polar so has dipole forces, unlike the other nonpolar covalent compounds. 342 CHAPTER 10 f. LIQUIDS AND SOLIDS I2; I2 has only LD forces while CsBr and CaO have much stronger ionic forces. I2 has weakest intermolecular forces so has smallest ΔHfusion. Properties of Liquids 37. The attraction of H2O for glass is stronger than the H2O‒H2O attraction. The miniscus is concave to increase the area of contact between glass and H2O. The Hg‒Hg attraction is greater than the Hg‒glass attraction. The miniscus is convex to minimize the Hg‒glass contact. 38. A molecule at the surface of a waterdrop is subject to attractions only by molecules below it and to each side. The effect of this uneven pull on the surface molecules tends to draw them into the body of the liquid and causes the droplet to assume the shape that has the minimum surface area, a sphere. 39. The structure of H2O2 is H‒O‒O‒H, which produces greater hydrogen bonding than water. Long chains of hydrogen‒bonded H2O2 molecules then get tangled together. 40. CO2 is a gas at room temperature. As mp and bp increase, the strength of the intermolecular forces also increases. Therefore, the strength of forces is CO2 < CS2 < CSe2. From a structural standpoint this is expected. All three are linear, nonpolar molecules. Thus, only London dispersion forces are present. Since the molecules increase in size from CO2 < CS2 < CSe2, the strength of the intermolecular forces will increase in the same order. Structures and Properties of Solids 41. nλ = 2d sin θ, d = 42. d= 43. = 44. nλ = 2d sin θ, d = nλ 2 154 pm = = 408 pm = 4.08 × 10 10 m 2 sin θ 2 sin 22.20 2 1.36 10 10 m sin 15.0 o 2 d sin θ = = 7.04 × 10 11 m = 0.704 1 n sin θ = 45. nλ 1 154 pm = = 313 pm = 3.13 × 10 10 m 2 sin θ 2 sin 14.22 1 2.63 nλ = = 4.91 = 4.91 × 10 10 m = 491 pm 2 sin 15.55 2 sin θ 2 2.63 nλ = = 0.536, θ = 32.4° 2 4.91 2d A cubic closest packed structure has a face-centered cubic unit cell. In a face-centered cubic unit, there are: 8 corners × 1/ 8 atom 1 / 2 atom = 4 atoms 6 faces corner face CHAPTER 10 LIQUIDS AND SOLIDS 343 The atoms in a face-centered cubic unit cell touch along the face diagonal of the cubic unit cell. Using the Pythagorean formula where l = length of the face diagonal and r = radius of the atom: l2 + l2 = (4r)2 4r l 2 l2 = 16 r2 l=r 8 l l=r 8 = 197 × 10 12 m × 8 = 5.57 × 10 10 m = 5.57 × 10 8 cm Volume of a unit cell = l 3 = (5.57 × 10 8 cm)3 = 1.73 × 10 22 cm3 Mass of a unit cell = 4 Ca atoms × density = 46. 1 mol Ca 40.08 g Ca = 2.662 × 10 22 g Ca 23 mol Ca 6.022 10 atoms mass 2.662 10 22 g = 1.54 g/cm3 22 3 volume 1.73 10 cm There are 4 Ni atoms in each unit cell: For a unit cell: density = mass = 6.84 g/cm3 = volume 4 Ni atoms 1 mol Ni 58.69 g Ni 23 mol Ni 6.022 10 atoms 3 l Solving: l = 3.85 × 10 8 cm = cube edge length For a face centered cube: (4r)2 = l 2 + l 2 = 2 l 2 r 8 = l, r = l / 8 4r l r = 3.85 × 10 8 cm / 8 r = 1.36 × 10 8 cm = 136 pm l 344 47. CHAPTER 10 LIQUIDS AND SOLIDS The unit cell for cubic closest packing is the face-centered unit cell. The volume of a unit cell is: V = l 3 = (492 × 10 10 cm)3 = 1.19 × 10 22 cm3 There are 4 Pb atoms in the unit cell, as is the case for all face-centered cubic unit cells. The mass of atoms in a unit cell is: mass = 4 Pb atoms × density = 1 mol Pb 207.2 g Pb = 1.38 × 10 21 g 23 mol Pb 6.022 10 atoms mass 1.38 10 21 g = 11.6 g/cm3 volume 1.19 10 22 cm3 From Exercise 45, the relationship between the cube edge length, l, and the radius of an atom in a face-centered unit cell is: l = r 8 . r= l = 8 48. 492 pm 8 = 174 pm = 1.74 × 10 10 m A face-centered cubic unit cell contains 4 atoms. For a unit cell: mass of X = volume × density = (4.09 × 10 8 cm)3 × 10.5 g/cm3 = 7.18 × 10 22 g mol X = 4 atoms X × Molar mass = 49. 1 mol X = 6.642 × 10 24 mol X 6.022 10 23 atoms 7.18 10 22 g X = 108 g/mol; The metal is silver (Ag). 6.642 10 24 mol X For a body-centered unit cell: 8 corners × 1 / 8 Ti + Ti at body center = 2 Ti atoms corner All body-centered unit cells have 2 atoms per unit cell. For a unit cell: 2 atoms Ti density = 4.50 g/cm3 = 1 mol Ti 47.88 g Ti mol Ti 6.022 10 23 atoms 3 l Solving: l = edge length of unit cell = 3.28 × 10 8 cm = 328 pm Assume Ti atoms just touch along the body diagonal of the cube, so body diagonal = 4 × radius of atoms = 4r. CHAPTER 10 LIQUIDS AND SOLIDS 345 The triangle we need to solve is: 4r l l 2 3.28 x 10 -8 cm (3.28 x 10 -8cm) 2 (4r)2 = (3.28 × 10 8 cm)2 + [(3.28 × 10 8 cm) 2 ]2, r = 1.42 × 10 8 cm = 142 pm For a body-centered unit cell (bcc), the radius of the atom is related to the cube edge length by: 4r = l 3 or l = 4r/ 3 . 50. From Exercise 10.49: 4r 16 r2 = l 2 + 2 l 2 l = 4r/ 3 = 2.309 r l l = 2.309 (222 pm) = 513 pm = 5.13 × 10 8 cm l 2 In a bcc, there are 2 atoms/unit cell. For a unit cell: density = 51. mass volume 2 atoms Ba 1 mol Ba 137.3 g Ba 23 mol Ba 3.38 g 6.022 10 atoms 8 3 (5.13 10 cm) cm3 If a face-centered cubic structure, then 4 atoms/unit cell and from Exercise 10.45: 2 l 2 = 16 r2 4r l l = r 8 = (144 pm) 8 = 407 pm l = 4.07 × 10 10 m = 4.07 × 10 8 cm l 4 atoms Au density = 1 mol Au 197 .0 g Au 23 mol Au 6.022 10 atoms = 19.4 g/cm3 8 3 (4.07 10 cm) 346 CHAPTER 10 LIQUIDS AND SOLIDS If a body-centered cubic structure, then 2 atoms/unit cell and from Exercise 10.49: 16 r2 = l 2 + 2 l 2 4r l l = 4r/ 3 = 333 pm = 3.33 × 10 10 m l = 3.33 × 10 8 cm l 2 2 atoms Au density = 1 mol Au 197 .0 g Au 23 mol Au 6.022 10 atoms = 17.7 g/cm3 8 3 (3.33 10 cm) The measured density is consistent with a face-centered cubic unit cell. 52. If face-centered cubic: l = r 8 = (137 pm) 8 = 387 pm = 3.87 × 10 8 cm 4 atoms W density = 1 mol 183 .9 g W 23 mol 6.022 10 atoms = 21.1 g/cm3 8 3 (3.87 10 cm) If body-centered cubic: l = 4r 3 = 4 137 pm 3 2 atoms W density = = 316 pm = 3.16 × 10 8 cm 1 mol 183 .9 g W mol 6.022 10 23 atoms = 19.4 g/cm3 8 3 (3.16 10 cm) The measured density is consistent with a body-centered unit cell. 53. In a face-centered unit cell (ccp structure), the atoms touch along the face diagonal: (4r)2 = l 2 + l2 4r l l=r 8 Vcube = l 3 = ( r 8 )3 = 22.63 r3 There are four atoms in a face-centered cubic cell (see Exercise 10.45). Each atom has a volume of 4/3 πr3. CHAPTER 10 LIQUIDS AND SOLIDS Vatoms = 4 × 4 3 347 πr3 = 16.76 r3 Vatoms 16 .76 r 3 = 0.7406 or 74.06% of the volume of each unit cell is occupied by Vcube 22 .63 r 3 atoms. So, In a simple cubic unit cell, the atoms touch along the cube edge (l): 2(radius) = 2r = l 2r l Vcube = l3 = (2r)3 = 8 r3 There is one atom per simple cubic cell (8 corner atoms × 1/8 atom per corner = 1 atom/unit cell). Each atom has an assumed volume of 4/3 πr3 = volume of a sphere. Vatom = 4 3 πr3 = 4.189 r3 Vatom 4.189 r 3 = = 0.5236 or 52.36% of the volume of each unit cell is occupied by Vcube 8 r3 atoms. So, A cubic closest packed structure packs the atoms much more efficiently than a simple cubic structure. 54. From Exercise 10.49, a body-centered unit cell contains 2 net atoms, and the length of a cube edge (l) is related to the radius of the atom (r) by the equation l = 4r/ 3 . Volume of unit cell = l3 = (4 r/ 3 )3 = 12.32 r3 Volume of atoms in unit cell = 2 × So: 4 3 π r3 = 8.378 r3 Vatoms 8.378 r 3 = 0.6800 = 68.00% occupied Vcube 12 .32 r 3 To determine the radius of the Fe atoms, we need to determine the cube edge length (l). 1 mol Fe 55.85 g Fe 1 cm3 Volume of unit cell = 2 Fe atoms mol Fe 7.86 g 6.022 10 23 atoms = 2.36 × 10 23 cm3 Volume = l 3 = 2.36 × 10 23 cm3, l = 2.87 × 10 8 cm l = 4r/ 3 , r = l 3 / 4 = 2.87 × 10 8 cm × 3 / 4 = 1.24 × 10 8 cm 348 CHAPTER 10 LIQUIDS AND SOLIDS 55. Doping silicon with phosphorus produces an n-type semiconductor. The phosphorus adds electrons at energies near the conduction band of silicon. Electrons do not need as much energy to move from filled to unfilled energy levels so conduction increases. Doping silicon with gallium produces a p-type semiconductor. Because gallium has fewer valence electrons than silicon, holes (unfilled energy levels) at energies in the previously filled molecular orbitals are created, which induces greater electron movement (greater conductivity). 56. A rectifier is a device that produces a current that flows in one direction from an alternating current which flows in both directions. In a p-n junction, a p-type and an n-type semiconductor are connected. The natural flow of electrons in a p-n junction is for the excess electrons in the n-type semiconductor to move to the empty energy levels (holes) of the ptype semiconductor. Only when an external electric potential is connected so that electrons flow in this natural direction will the current flow easily (forward bias). If the external electric potential is connected in reverse of the natural flow of electrons, no current flows through the system (reverse bias). A p-n junction only transmits a current under forward bias, thus converting the alternating current to direct current. 57. In has fewer valence electrons than Se, thus, Se doped with In would be a p-type semiconductor. 58. To make a p-type semiconductor we need to dope the material with atoms that have fewer valence electrons. The average number of valence electrons is four when 50-50 mixtures of group 3A and group 5A elements are considered. We could dope with more of the Group 3A element or with atoms of Zn or Cd. Cadmium is the most common impurity used to produce p-type GaAs semiconductors. To make an n-type GaAs semiconductor, dope with an excess group 5A element or dope with a Group 6A element such as sulfur. 59. Egap = 2.5 eV × 1.6 × 10 19 J/eV = 4.0 × 10 19 J; We want Egap = Elight, so: λ hc (6.63 10 34 J s) (3.00 10 8 m / s) = = 5.0 × 10 7 m = 5.0 × 102 nm 4.0 10 19 J E hc (6.63 10 34 J s) (2.998 10 8 m / s) = 2.72 × 10 19 J = energy of band gap 9 λ 730 . 10 m 60. E= 61. a. 8 corners × 1/ 8 Cl 1 / 2 Cl + 6 faces × = 4 Cl ions face corner 1/ 4 Na + 1 Na at body center = 4 Na ions; NaCl is the formula. edge 1/ 8 Cl b. 1 Cs ion at body center; 8 corners × = 1 Cl ion; CsCl is the formula. corner 12 edges × c. There are 4 Zn ions inside the cube. 8 corners × 1/ 8 S 1/ 2 S + 6 faces × = 4 S ions; ZnS is the formula. face corner CHAPTER 10 LIQUIDS AND SOLIDS d. 8 corners × 4 faces × 62. 349 1/ 8 Ti + 1 Ti at body center = 2 Ti ions corner 1/ 2 O + 2 O inside cube = 4 O ions; TiO2 is the formula. face Both As ions are inside the unit cell. 8 corners × 1/ 8 Ni 1 / 4 Ni + 4 edges × = 2 Ni ions corner edge The unit cell contains 2 ions of Ni and 2 ions of As which gives a formula of NiAs. 63. There is one octahedral hole per closest packed anion in a closest packed structure. If half of the octahedral holes are filled, there is a 2:1 ratio of fluoride ions to cobalt ions in the crystal. The formula is CoF2. 64. There are 2 tetrahedral holes per closest packed anion. Let f = fraction of tetrahedral holes filled by the cations. Na2O: cation to anion ratio = 2 2f = , f = 1; All of the tetrahedral holes are filled by Na+ 1 1 cations. CdS: cation to anion ratio = 1 2f 1 = , f = ; 1/2 of the tetrahedral holes are filled by 1 1 2 Cd2+ cations. ZrI4: cation to anion ratio = 1 2f 1 = , f = ; 1/8 of the tetrahedral holes are filled by 4 1 8 Zr4+ cations. 65. In a cubic closest packed array of anions, there are twice the number of tetrahedral holes as anions present and an equal number of octahedral holes as anions present. A cubic closest packed array of sulfide ions will have 4 S2 ions, 8 tetrahedral holes, and 4 octahedral holes. In this structure we have 1/8(8) = 1 Zn2+ ion and 1/2(4) = 2 Al3+ ions present along with the 4 S2 ions. The formula is ZnAl2S4. 66. The two-dimension unit cell is: Assuming the anions A are the larger circles, there are 4 anions completely in the unit cell. The corner cations (smaller circles) are shared by 4 different unit cells. Therefore, there is 1 cation in the middle of the unit cell plus 1/4(4) = 1 net cation from the corners. Each unit cell has 2 cations and 4 anions. The empirical formula is MA2. 350 CHAPTER 10 LIQUIDS AND SOLIDS 67. 8 F ions at corners × 1/8 F/corner = 1 F ion per unit cell; Because there is one cubic hole per cubic unit cell, there is a 2:1 ratio of F ions to metal ions in the crystal if only ½ of the body centers are filled with the metal ions. The formula is MF2 where M2+ is the metal ion. 68. Mn ions at 8 corners: 8(1/8) = 1 Mn ion; F ions at 12 edges: 12(1/4) = 3 F ions Formula is MnF3. Assuming fluoride is -1 charged, the charge on Mn is +3. 69. From Fig. 10.37, MgO has the NaCl structure containing 4 Mg2+ ions and 4 O2 ions per facecentered unit cell. 4 MgO formula units × 1 mol MgO 40.31 g MgO = 2.678 × 10 22 g MgO 23 1 mol MgO 6.022 10 atoms Volume of unit cell = 2.678 × 10 22 g MgO × 1 cm3 = 7.48 × 10 23 cm3 3.58 g Volume of unit cell = l3, l = cube edge length; l = (7.48 × 10 23 cm3)1/3 = 4.21 × 10 8 cm For a face-centered unit cell, the O2 ions touch along the face diagonal: 2 l 4 rO 2 , rO 2 2 4.21 10 8 cm = 1.49 × 10 8 cm 4 The cube edge length goes through two radii of the O2 anions and the diameter of the Mg2+ cation. So: l = 2 rO 2 2 rMg 2 , 4.21 × 10 8 cm = 2(1.49 × 10 8 cm) + 2 rMg 2 , rMg 2 = 6.15 × 10 9 cm 70. CsCl is a simple cubic array of Cl ions with Cs+ in the middle of each unit cell. There is one Cs+ and one Cl ion in each unit cell. Cs+ and Cl ions touch along the body diagonal. body diagonal = 2 rCs 2 rCl 3 l , l = length of cube edge In each unit cell: 168 .4 g CsCl 1 mol CsCl mass = 1 CsCl formula unit 23 6.022 10 formula units mol CsCl = 2.796 × 10 22 g 1 cm3 volume = l 3 = 2.796 × 10 22 g CsCl × = 7.04 × 10 23 cm3 3.97 g CsCl l 3 = 7.04 × 10 23 cm3, l = 4.13 × 10 8 cm = 413 pm = length of cube edge 2 rCs 2 rCl 3 l 3(413 pm) = 715 pm CHAPTER 10 LIQUIDS AND SOLIDS 351 The distance between ion centers = rCs rCl = 715 pm/2 = 358 pm From ionic radii: rCs = 169 pm and rCl = 181 pm; rCs rCl = 169 + 181 = 350. pm The actual distance is 8 pm (2.3%) greater than that calculated from values of ionic radii. 71. 72. 73. a. CO2: molecular b. SiO2: network c. Si: atomic, network d. CH4: molecular e. Ru: atomic, metallic f. I2: molecular g. KBr: ionic h. H2O: molecular i. NaOH: ionic j. k. CaCO3: ionic l. PH3: molecular U: atomic, metallic a. diamond: atomic, network b. PH3: molecular c. H2: molecular d. Mg: atomic, metallic e. KCl: ionic f. quartz: network g. NH4NO3: ionic h. SF2: molecular i. Ar: atomic, group 8A j. k. C6H12O6: molecular Cu: atomic, metallic a. The unit cell consists of Ni at the cube corners and Ti at the body center, or Ti at the cube corners and Ni at the body center. b. 8 × 1/8 = 1 atom from corners + 1 atom at body center; Empirical formula = NiTi c. Both have a coordination number of 8 (both are surrounded by 8 atoms). 74. 8 corners × 1 / 8 Xe 1/ 4 F + 1 Xe inside cell = 2 Xe; 8 edges × + 2 F inside cell = 4 F edge corner Empirical formula is XeF2. This is also the molecular formula. 75. Structure 1 8 corners × 6 faces × 1 / 8 Ca = 1 Ca atom corner 1/ 2 O = 3 O atoms face 1 Ti at body center. Formula = CaTiO3 Structure 2 8 corners × 1/ 8 Ti = 1 Ti atom corner 12 edges × 1/ 4 O = 3 O atoms corner 1 Ca at body center. Formula = CaTiO3 In the extended lattice of both structures, each Ti atom is surrounded by six O atoms. 76. With a cubic closest packed array of oxygen ions, we have 4 O2 ions per unit cell. We need to balance the total 8 charge of the anions with a +8 charge from the Al3+ and Mg2+ cations. 352 CHAPTER 10 LIQUIDS AND SOLIDS The only combination of ions that gives a +8 charge is 2 Al3+ ions and 1 Mg2+ ion. The formula is Al2MgO4. There are an equal number of octahedral holes as anions (4) in a cubic closest packed array, and twice the number of tetrahedral holes as anions in a cubic closest packed array. For the stoichiometry to work out, we need 2 Al3+ and 1 Mg2+ per unit cell. Hence, one-half of the octahedral holes are filled with Al3+ ions and one-eighth of the tetrahedral holes are filled with Mg2+ ions. 77. a. Y: 1 Y in center; Ba: 2 Ba in center Cu: 8 corners × O: 20 edges × 1 / 8 Cu 1 / 4 Cu = 1 Cu, 8 edges × = 2 Cu, total = 3 Cu atoms corner edge 1/ 2 O 1/ 4 O = 5 oxygen, 8 faces × = 4 oxygen, total = 9 O atoms face edge Formula: YBa2Cu3O9 b. The structure of this superconductor material follows the second perovskite structure described in Exercise 10.75. The YBa2Cu3O9 structure is three of these cubic perovskite unit cells stacked on top of each other. The oxygen atoms are in the same places, Cu takes the place of Ti, two of the calcium atoms are replaced by two barium atoms, and one Ca is replaced by Y. c. Y, Ba, and Cu are the same. Some oxygen atoms are missing. 12 edges × 1/ 2 O 1/ 4 O = 3 O, 8 faces × = 4 O, total = 7 O atoms face edge Superconductor formula is YBa2Cu3O7. 78. a. Structure (a): 1/ 8 Tl = 1 Tl corner Ba: 2 Ba inside unit cell; Tl: 8 corners × 1 / 4 Cu = 1 Cu edge 1/ 2 O 1/ 4 O O: 6 faces × + 8 edges × = 5 O; Formula = TlBa2CuO5 face edge Cu: 4 edges × Structure (b): Tl and Ba are the same as in structure (a). Ca: 1 Ca inside unit cell; Cu: 8 edges × O: 10 faces × 1 / 4 Cu = 2 Cu edge 1/ 2 O 1/ 4 O + 8 edges × = 7 O; Formula = TlBa2CaCu2O7 face edge CHAPTER 10 LIQUIDS AND SOLIDS 353 Structure (c): Tl and Ba are the same, and two Ca are located inside the unit cell. Cu: 12 edges × 1/ 2 O 1 / 4 Cu 1/ 4 O = 3 Cu; O: 14 faces × + 8 edges × =9O edge face edge Formula: TlBa2Ca2Cu3O9 Structure (d): Following similar calculations, formula = TlBa2Ca3Cu4O11 b. Structure (a) has one planar sheet of Cu and O atoms, and the number increases by one for each of the remaining structures. The order of superconductivity temperature from lowest to highest temperature is: (a) < (b) < (c) < (d). c. TlBa2CuO5: 3 + 2(2) + x + 5(2) = 0, x = +3 Only Cu3+ is present in each formula unit. TlBa2CaCu2O7: 3 + 2(2) + 2 + 2(x) + 7(2) = 0, x = +5/2 Each formula unit contains 1 Cu2+ and 1 Cu3+. TlBa2Ca2Cu3O9: 3 + 2(2) + 2(2) + 3(x) + 9(2) = 0, x = +7/3 Each formula unit contains 2 Cu2+ and 1 Cu3+. TlBa2Ca3Cu4O11: 3 + 2(2) + 3(2) + 4(x) + 11(2) = 0, x = +9/4 Each formula unit contains 3 Cu2+ and 1 Cu3+. d. This superconductor material achieves variable copper oxidation states by varying the numbers of Ca, Cu and O in each unit cell. The mixtures of copper oxidation states are discussed above. The superconductor material in Exercise 10.77 achieves variable copper oxidation states by omitting oxygen at various sites in the lattice. Phase Changes and Phase Diagrams 79. If we graph ln Pvap vs 1/T, the slope of the resulting straight line will be ΔHvap/R. Pvap ln Pvap T (Li) 1 torr 10. 100. 400. 760. 0 2.3 4.61 5.99 6.63 1023 K 1163 1353 1513 1583 1/T 9.775 × 104 K1 8.598 × 104 7.391 × 104 6.609 × 104 6.317 × 104 T (Mg) 893 K 1013 1173 1313 1383 1/T 11.2 × 104 K1 9.872 × 104 8.525 × 104 7.616 × 104 7.231 × 104 354 CHAPTER 10 LIQUIDS AND SOLIDS For Li: We get the slope by taking two points (x, y) that are on the line we draw. For a line: slope = y y Δy 2 1 Δx x 2 x1 or we can determine the straight line equation using a calculator. The general straight line equation is y = mx + b where m = slope and b = y-intercept. The equation of the Li line is: ln Pvap = 1.90 × 104(1/T) + 18.6, slope = 1.90 × 104 K Slope = ΔHvap/R, ΔHvap = slope × R = 1.90 × 104 K × 8.3145 J/Kmol ΔHvap = 1.58 × 105 J/mol = 158 kJ/mol For Mg: The equation of the line is: ln Pvap = 1.67 × 104(1/T) + 18.7, slope = 1.67 × 104 K ΔHvap = -slope × R = 1.67 × 104 K × 8.3145 J/Kmol, ΔHvap = 1.39 × 105 J/mol = 139 kJ/mol The bonding is stronger in Li since ΔHvap is larger for Li. 80. We graph ln Pvap vs 1/T. The slope of the line equals ΔHvap/R. CHAPTER 10 LIQUIDS AND SOLIDS 103/T (K1) Pvap (torr) T(K) 273 283 293 303 313 323 353 slope = 3.66 3.53 3.41 3.30 3.19 3.10 2.83 14.4 26.6 47.9 81.3 133 208 670. 355 ln Pvap 2.67 3.28 3.87 4.40 4.89 5.34 6.51 6.6 2.5 (2.80 10 -4600 K = 3 ΔH vap R = 4600 K 3.70 10 3 ) K 1 ΔH vap 8.3145 J / K mol , ΔHvap = 38,000 J/mol = 38 kJ/mol To determine the normal boiling point, we can use the following formula: P ΔH vap 1 1 = ln 1 = R P2 T2 T1 At the normal boiling point, the vapor pressure equals 1.00 atm or 760. torr. At 273 K, the vapor pressure is 14.4. torr (from data in the problem). 38,000 J / mol 1 1 14.4 , 3.97 = 4.6 × 103 (1/T2 3.66 × 103) ln = 760 . 8.3145 J / K mol T2 273 K 8.6 × 104 + 3.66 × 103 = 1/T2 = 2.80 × 103, T2 = 357 K = normal boiling point 81. At 100.°C (373 K), the vapor pressure of H2O is 1.00 atm = 760. torr. For water, ΔHvap = 40.7 kJ/mol. P ln 1 P2 ΔH vap 1 P 1 or ln 2 P R T2 T1 1 ΔH vap 1 1 R T1 T2 520 . torr 40.7 10 3 J / mol 1 1 , 7.75 × 105 = = ln 760 . torr 8 . 3145 J / K mol 373 K T 2 7.75 × 105 = 2.68 × 103 82. 1 1 373 K T2 1 1 1 = 2.76 × 103, T2 = = 362 K or 89°C , 2.76 10 3 T2 T2 40.7 10 3 J / mol 1 1 P , ln P2 = 5.27, P2 = e5.27 = 194 atm ln 2 = 8 . 3145 J / K mol 373 K 623 K 1 . 00 356 83. CHAPTER 10 P ln 1 P2 LIQUIDS AND SOLIDS ΔH vap ΔH vap 1 836 torr 1 1 1 , ln R T2 T1 213 torr 8.3145 J / K mol 313 K 353 K Solving: Hvap = 3.1 × 104 J/mol; For the normal boiling point, P = 1.00 atm = 760. torr. 760 . torr 3.1 10 4 J / mol 1 1 1 1 = = 3.4 × 10 4 ln , 213 torr 8 . 3145 J / K mol 313 K T 313 T 1 1 T1 = 350. K = 77C; The normal boiling point of CCl4 is 77C. 84. P ln 1 P2 ΔH vap 1 1 R T2 T1 P1 = 760. torr, T1 = 56.5°C + 273.2 = 329.7 K; P2 = 630. torr, T2 = ? 760 . ln 630 . 32.0 10 3 J / mol 1 1 = 8.3145 J / K mol T2 329 .7 , 0.188 = 3.85 × 103 1 3.033 10 3 T2 1 1 3 5 3.033 × 10 = 4.88 × 10 , = 3.082 × 10 3 , T2 = 324.5 K = 51.3°C T2 T2 630 . torr ln P2 32.0 10 3 J / mol 1 1 = 8.3145 J / K mol 298.2 324.5 ln P2 = 5.40, P2 = e5.40 = 221 torr 85. , ln 630. ln P2 = 1.05 CHAPTER 10 86. LIQUIDS AND SOLIDS 357 X(g, 100.°C) → X(g, 75°C), ΔT = -25°C q1 = sgas × m × ΔT = 1.0 J × 250. g × (25°C) = 6300 J = 6.3 kJ gC X(g, 75°C) → X(l, 75°C), q2 = 250. g × X(l, 75°C) → X(l, -15°C), q3 = 1 mol 20. kJ = 67 kJ 75.0 g mol 2.5 J × 250. g × (-90.°C) = -56,000 J = 56 kJ gC X(l, 15°C) → X(s, 15°C), q4 = 250. g × X(s, 15°C) → X(s, 50.°C), q5 = 1 mol 5.0 kJ = 17 kJ 75.0 g mol 3.0 J × 250. g × (35°C) = 26,000 J = 26 kJ gC qtotal = q1 + q2 + q3 + q4 + q5 = -6.3 67 - 56 17 26 = 172 kJ 87. H2O(s, 20.°C) → H2O(s, 0°C), ΔT = 20.°C q1 = sice × m × ΔT = 2.03 J × 5.00 × 102 g × 20.°C = 2.0 × 104 J = 20. kJ gC H2O(s, 0°C) → H2O(l, 0°C), q2 = 5.00 × 102 g H2O × H2O(l, 0°C) → H2O(l, 100.°C), q3 = 1 mol 6.02 kJ = 167 kJ 18.02 g mol 4.2 J × 5.00 × 102 g × 100.°C = 2.1 × 105 J = 210 kJ gC H2O(l, 100.°C) → H2O(g, 100.°C), q4 = 5.00 × 102 g × H2O(g, 100.°C) → H2O(g, 250.°C), q5 = 1 mol 40.7 kJ = 1130 kJ 18.02 g mol 2.0 J × 5.00 × 102 g × 150.°C = 1.5 × 105 J gC = 150 kJ qtotal = q1 + q2 + q3 + q4 + q5 = 20. + 167 + 210 + 1130 + 150 = 1680 kJ 88. H2O(g, 125°C) → H2O(g, 100.°C), q1 = 2.0 J/g °C × 75.0 g × (25°C) = -3800 J = 3.8 kJ H2O(g, 100.°C) → H2O(l, 100.°C), q2 = 75.0 g × 1 mol 40.7 kJ = 169 kJ 18.02 g mol H2O(l, 100.°C) → H2O(l, 0°C), q3 = 4.2 J/g °C × 75.0 g × (-100.°C) = 32,000 J = 32 kJ To convert H2O(g) at 125°C to H2O(l) at 0°C requires (3.8 kJ 169 kJ 32 kJ =) 205 kJ of heat removed. To convert from H2O(l) at 0°C to H2O(s) at 0°C requires: 358 CHAPTER 10 LIQUIDS AND SOLIDS 1 mol 6.02 kJ = 25 kJ 18.02 g mol This amount of energy puts us over the -215 kJ limit (-205 kJ - 25 kJ = -230. kJ). Therefore, a mixture of H2O(s) and H2O(l) will be present at 0°C when 215 kJ of heat are removed from the gas sample. q4 = 75.0 g × 89. Total mass H2O = 18 cubes × 30.0 g 1 mol H 2O = 540. g; 540. g H2O × = 30.0 mol H2O cube 18.02 g Heat removed to produce ice at -5.0°C: 6.02 10 3 J 4.18 J 2.03 J × 540. g × 22.0 °C + × 30.0 mol + × 540. g × 5.0 °C mol gC gC 4.97 × 104 J + 1.81 × 105 J + 5.5 × 103 J = 2.36 × 105 J 1 g CF2Cl2 = 1.49 × 103 g CF2Cl2 must be vaporized. 158 J 1 mol 368 kJ Heat released = 0.250 g Na × = 2.00 kJ 22.99 g 2 mol 2.36 × 105 J × 90. To melt 50.0 g of ice requires: 50.0 g ice × 1 mol H 2 O 6.02 kJ = 16.7 kJ 18.02 g mol The reaction doesn't release enough heat to melt all of the ice. The temperature will remain at 0°C. 91. A: solid B: liquid D: solid + vapor E: solid + liquid + vapor F: liquid + vapor G: liquid + vapor triple point: E C: vapor H: vapor critical point: G normal freezing point: temperature at which solid - liquid line is at 1.0 atm (see plot below). normal boiling point: temperature at which liquid - vapor line is at 1.0 atm (see plot below). 1.0 atm nfp nbp Since the solid-liquid line has a positive slope, the solid phase is denser than the liquid phase. CHAPTER 10 92. LIQUIDS AND SOLIDS 359 a. 3 b. Triple point at 95.31°C: rhombic, monoclinic, gas Triple point at 115.18°C: monoclinic, liquid, gas Triple point at 153°C: rhombic, monoclinic, liquid c. From the phase diagram, the monoclinic solid phase is stable at T = 100°C and P = 1 atm. d. Normal melting point = 115.21°C; normal boiling point = 444.6°C; The normal melting and boiling points occur at P = 1.0 atm. e. Rhombic is the densest phase since the rhombic-monoclinic equilibrium line has a positive slope and since the solid-liquid lines also have positive slopes. No; P = 1.0 × 10 5 atm is at a pressure somewhere between the 95.31°C and 115.18°C triple points. At this pressure, the rhombic and gas phases are never in equilibrium with each other, so rhombic sulfur cannot sublime at P = 1.0 × 10 5 atm. However, monoclinic sulfur can sublime at this pressure. g. From the phase diagram, we would start off with gaseous sulfur. At 100°C and ~1 × 10 5 atm, S(g) would convert to the solid monoclinic form of sulfur. Finally at 100°C and some large pressure less than 1420 atm, S(s, monoclinic) would convert to the solid rhombic form of sulfur. Summarizing, the phase changes are S(g) ➔ S(monoclinic) ➔ S(rhombic). f. 93. a. two b. Higher pressure triple point: graphite, diamond and liquid; Lower pressure triple point: graphite, liquid and vapor c. It is converted to diamond (the more dense solid form). d. Diamond is more dense, which is why graphite can be converted to diamond by applying pressure. 94. The following sketch of the Br2 phase diagram is not to scale. Since the triple point of Br2 is at a temperature below the freezing point of Br2, the slope of the solid-liquid line is positive. 100 Solid Pressure (atm) Liquid 1 0.05 Gas -7.3 -7.2 59 Temperature ( oC) 320 360 CHAPTER 10 LIQUIDS AND SOLIDS The positive slopes of all the lines indicate that Br2(s) is more dense than Br2(l) which is more dense than Br2(g). At room temperature (~22°C) and 1 atm, Br 2(l) is the stable phase. Br2(l) cannot exist at a temperature below the triple point temperature of 7.3°C and at a temperature above the critical point temperature of 320°C. The phase changes that occur as temperature is increased at 0.10 atm are solid ➔ liquid ➔ gas. 95. Because the density of the liquid phase is greater than the density of the solid phase, the slope of the solid-liquid boundary line is negative (as in H2O). With a negative slope, the melting points increase with a decrease in pressure so the normal melting point of X should be greater than 225C. 96. 760 s l P (torr) g 280 -121 -112 -107 T(C) From the three points given, the slope of the s-l boundary line is positive so Xe(s) is more dense than Xe(l). Also, the positive slope of this line tells us that the melting point of Xe increases as pressure increases. The same direct relationship exists for the boiling point of Xe as the l-g boundary line also has a positive slope. Additional Exercises 97. Chalk is composed of the ionic compound calcium carbonate (CaCO3). The electrostatic forces in ionic compounds are much stronger than the intermolecular forces in covalent compounds. Therefore, CaCO3 should have a much higher boiling point than the covalent compounds found in motor oil and in H2O. Motor oil is composed of nonpolar CC and CH bonds. The intermolecular forces in motor oil are therefore London dispersion forces. We generally consider these forces to be weak. However, with compounds that have large molar masses, these London dispersion forces add up significantly and can overtake the relatively strong hydrogen bonding interactions in water. CHAPTER 10 98. LIQUIDS AND SOLIDS Benzene 361 Naphthalene H H H H H H H H H H H H H LD forces only H LD forces only Note: London dispersion forces in molecules like benzene and naphthalene are fairly large. The molecules are flat, and there is efficient surface area contact among molecules. Large surface area contact leads to stronger London dispersion forces. Cl Carbon tetrachloride (CCl4) has polar bonds but is a nonpolar molecule. CCl4 only has LD forces. C Cl Cl Cl In terms of size and shape: CCl4 < C6H6 < C10H8 The strengths of the LD forces are proportional to size and are related to shape. Although CCl4 is fairly large, its overall spherical shape gives rise to relatively weak LD forces as compared to flat molecules like benzene and naphthalene. The physical properties given in the problem are consistent with the order listed above. Each of the physical properties will increase with an increase in intermolecular forces. Acetone Acetic Acid O O H 3C C CH 3 LD, dipole H 3C C O H LD, dipole, H bonding 362 CHAPTER 10 LIQUIDS AND SOLIDS Benzoic acid O C O H LD, dipole, H bonding We would predict the strength of intermolecular forces for the last three molecules to be: acetone < acetic acid < benzoic acid polar H bonding H bonding, but large LD forces because of greater size and shape. This ordering is consistent with the values given for bp, mp, and ΔHvap. The overall order of the strengths of intermolecular forces based on physical properties are: acetone < CCl4 < C6H6 < acetic acid < naphthalene < benzoic acid The order seems reasonable except for acetone and naphthalene. Since acetone is polar, we would not expect it to boil at the lowest temperature. However, in terms of size and shape, acetone is the smallest molecule, and the LD forces in acetone must be very small compared to the other molecules. Naphthalene must have very strong LD forces because of its size and flat shape. 99. At any temperature, the plot tells us that substance A has a higher vapor pressure than substance B, with substance C having the lowest vapor pressure. Therefore, the substance with the weakest intermolecular forces is A, and the substance with the strongest intermolecular forces is C. NH3 can form hydrogen bonding interactions while the others cannot. Substance C is NH3. The other two are nonpolar compounds with only London dispersion forces. Since CH4 is smaller than SiH4, CH4 will have weaker LD forces and is substance A. Therefore, substance B is SiH4. 100. As the electronegativity of the atoms covalently bonded to H increases, the strength of the hydrogen bonding interaction increases. N ⋅⋅⋅ H‒N < N ⋅⋅⋅ H‒ O < O ⋅⋅⋅ H‒O < O ⋅⋅⋅ H‒F < F ⋅⋅⋅ H‒F weakest 101. strongest If TiO2 conducts electricity as a liquid, then it is an ionic solid; if not, then TiO2 is a network solid. CHAPTER 10 LIQUIDS AND SOLIDS 363 102. One B atom and one N atom together have the same number of electrons as two C atoms. The description of physical properties sounds a lot like the properties of graphite and diamond, the two solid forms of carbon. The two forms of BN have structures similar to graphite and diamond. 103. B2H6: This compound contains only nonmetals so it is probably a molecular solid with covalent bonding. The low boiling point confirms this. 104. SiO2: This is the empirical formula for quartz, which is a network solid. CsI: This is a metal bonded to a nonmetal, which generally form ionic solids. The electrical conductivity in aqueous solution confirms this. W: Tungsten is a metallic solid as the conductivity data confirms. In order to set up an equation, we need to know what phase exists at the final temperature. To heat 20.0 g of ice from 10.0°C to 0.0°C requires: 2.03 J × 20.0 g × 10.0C = 406 J gC To convert ice to water at 0.0°C requires: q= q = 20.0 g × 1 mol 6.02 kJ = 6.68 kJ = 6680 J 18.02 mol To chill 100.0 g of water from 80.0°C to 0.0° requires: 4.18 J × 100.0 g × 80.0°C = 33,400 J of heat removed gC From the heat values above, the liquid phase exists once the final temperature is reached (a lot more heat is lost when the 100.0 g of water is cooled to 0.0°C than the heat required to convert the ice into water). To calculate the final temperature, we will equate the heat gain by the ice to the heat loss by the water. We will keep all quantities positive in order to avoid sign errors. The heat gain by the ice will be the 406 J required to convert the ice to 0.0°C plus the 6680 J required to convert the ice at 0.0°C into water at 0.0°C plus the heat required to raise the temperature from 0.0°C to the final temperature. q= heat gain by ice = 406 J + 6680 J + heat loss by water = 4.18 J × 20.0 g × (Tf 0.0°C) = 7.09 × 103 + 83.6 Tf gC 4.18 J × 100.0 g × (80.0°C Tf) = 3.34 × 104 418 Tf gC Solving for the final temperature: 7.09 × 103 + 83.6 Tf = 3.34 × 104 - 418 Tf, 502 Tf = 2.63 × 104, Tf = 52.4°C 105. 1.00 lb × 454 g 5 = 454 g H2O; A change of 1.00°F is equal to a change of C . lb 9 364 CHAPTER 10 The amount of heat in J in 1 Btu is: LIQUIDS AND SOLIDS 4.18 J 5 454 g C = 1.05 × 103 J = 1.05 kJ gC 9 It takes 40.7 kJ to vaporize 1 mol H2O (ΔHvap). Combining these: 1.00 10 4 Bu 1.05 kJ 1 mol H 2 O = 258 mol/hr hr Btu 40.7 kJ or: 106. 258 mol 18.02 g H 2O = 4650 g/hr = 4.65 kg/hr hr mol The critical temperature is the temperature above which the vapor cannot be liquefied no matter what pressure is applied. Since N2 has a critical temperature below room temperature (~22°C), it cannot be liquefied at room temperature. NH3, with a critical temperature above room temperature, can be liquefied at room temperature. Challenge Problems 107. H = qp = 30.79 kJ; E = qp + w, w = PV 101 .3 J L atm w = PV = 1.00 atm (28.90 L) = 28.9 L atm × = 2930 J E = 30.79 kJ + (2.93 kJ) = 27.86 kJ 108. XeCl2F2, 8 + 2(7) + 2(7) = 36 e Cl F F or Xe Cl Cl F Xe F polar (bond dipoles do not each other cancel) Cl nonpolar (bond dipoles cancel each other) These are two possible square planar molecular structures for XeCl2F2. One structure has the Cl atoms 90 apart, the other has the Cl atoms 180 apart. The structure with the Cl atoms 90 apart is polar; the other structure is nonpolar. The polar structure will have additional dipole forces so it has the stronger intermolecular forces and is the liquid. The gas form of XeCl2F2 is the nonpolar form having the Cl atoms 180 apart. 109. A single hydrogen bond in H2O has a strength of 21 kJ/mol. Each H2O molecule forms two H bonds. Thus, it should take 42 kJ/mol of energy to break all of the H bonds in water. Consider the phase transitions: solid 6.0 kJ liquid 40.7 kJ vapor Hsub = fusHvap CHAPTER 10 LIQUIDS AND SOLIDS 365 It takes a total of 46.7 kJ/mol to convert solid H2O to vapor (ΔHsub). This would be the amount of energy necessary to disrupt all of the intermolecular forces in ice. Thus, (42 ÷ 46.7) × 100 = 90% of the attraction in ice can be attributed to H bonding. 110. Both molecules are capable of H bonding. However, in oil of wintergreen the hydrogen bonding is intramolecular (within each molecule). O CH 3 C O H O In methyl-4-hydroxybenzoate, the H bonding is intermolecular, resulting in stronger forces between molecules and a higher melting point. 111. NaCl, MgCl2, NaF, MgF2 AlF3 all have very high melting points indicative of strong intermolecular forces. They are all ionic solids. SiCl4, SiF4, F2, Cl2, PF5 and SF6 are nonpolar covalent molecules. Only LD forces are present. PCl3 and SCl2 are polar molecules. LD forces and dipole forces are present. In these 8 molecular substances, the intermolecular forces are weak and the melting points low. AlCl3 doesn't seem to fit in as well. From the melting point, there are much stronger forces present than in the nonmetal halides, but they aren't as strong as we would expect for an ionic solid. AlCl3 illustrates a gradual transition from ionic to covalent bonding, from an ionic solid to discrete molecules. 112. a. The NaCl unit cell has a face centered cubic arrangement of the anions with cations in the octahedral holes. There are 4 NaCl formula units per unit cell and, since there is a 1:1 ratio of cations to anions in MnO, there would be 4 MnO formula units per unit cell, assuming an NaCl type structure. The CsCl unit cell has a simple cubic structure of anions with the cations in the cubic holes. There is one CsCl formula unit per unit cell, so there would be one MnO formula unit per unit cell if a CsCl structure is observed. molecules MnO 5.28 g MnO 1 mol MnO (4.47 10 8 cm)3 unit cell 70.94 g MnO cm3 6.022 10 23 molecules MnO = 4.00 molecules MnO mol MnO 4 From the calculation, MnO crystallizes in the NaCl type structure. b. From the NaCl structure and assuming the ions touch each other, then ℓ = cube edge length = 2rMn 2 2rO2 . ℓ = 4.47 × 10 8 cm = 2 rMn 2 + 2(1.40 × 10 8 cm), rMn 2 = 8.4 × 10 9 cm = 84 pm 366 113. CHAPTER 10 Out of 100.00 g: 28.31 g O × LIQUIDS AND SOLIDS 1 mol 1 mol = 1.769 mol O; 71.69 g Ti × = 1.497 mol Ti 16.00 g 47.88 g 1.497 1.769 = 1.182; = 0.8462; The formula is TiO1.182 or Ti0.8462O. 1.769 1.497 For Ti0.8462O, let x = Ti2+ per mol O2- and y = Ti3+ per mol O2-. Setting up two equations and solving: x + y = 0.8462 (mass balance) and 2x + 3y = 2 (charge balance); 2x + 3(0.8462 x) = 2 x = 0.539 mol Ti2+/mol O2- and y = 0.307 mol Ti3+/mol O2- 0.539 × 100 = 63.7% of the titanium ions are Ti2+ and 36.3% are Ti3+ (a 1.75:1 ion ratio). 0.8462 114. First we need to get the empirical formula of spinel. Assume 100.0 g of spinel. 37.9 g Al × 1 mol Al = 1.40 mol Al 26.98 g Al The mole ratios are 2:1:4. 17.1 g Mg × 1 mol Mg = 0.703 mol Mg 24.31 g Mg Empirical Formula = Al2MgO4 1 mol O 45.0 g O × = 2.81 mol O 16.00 g O Assume each unit cell contains an integral value (n) of Al2MgO4 formula units. Each Al2MgO4 formula unit has a mass of: 24.31 + 2(26.98) + 4(16.00) = 142.27 g/mol 1 mol 142 .27 g 23 mol 3.57 g 6.022 10 atoms , Solving: n = 8.00 8 3 (8.09 10 cm) cm3 n formula units density = Each unit cell has 8 formula units of Al2MgO4 or 16 Al3+, 8 Mg2+ and 32 O2 ions. 115. volume Cu density Mn massMn volume Cu massMn density Cu volume Mn massCu massCu volume Mn The type of cubic cell formed is not important; only that Cu and Mn crystallize in the same type of cubic unit cell is important. Each cubic unit cell has a specific relationship between the cube edge length, l, and the radius, r. In all cases l r. Therefore, V l3 r3. For the mass ratio, we can use the molar masses of Mn and Cu since each unit cell must contain the same number of Mn and Cu atoms. Solving: CHAPTER 10 LIQUIDS AND SOLIDS 367 volume Cu 54.94 g / mol (rCu ) 3 density Mn massMn density Cu massCu volume Mn 63.55 g / mol (1.056 rCu ) 3 3 density Mn 1 = 0.8645 × = 0.7341 density Cu 1.056 densityMn = 0.7341 × densityCu = 0.7341 × 8.96 g/cm3 = 6.58 g/cm3 116. a. The arrangement of the layers are: Layer 1 Layer 2 Layer 3 Layer 4 A total of 20 cannon balls will be needed. b. The layering alternates abcabc which is cubic closest packing. c. tetrahedron 117. a P c b As P is lowered, we go from a to b on the phase diagram. The water boils. The boiling of water is endothermic and the water is cooled (b →c), forming some ice. If the pump is left on, the ice will sublime until none is left. This is the basis of freeze drying. T 118. w = PΔV; V373 = Assume a constant P of 1.00 atm. nRT 1.00 (0.8206 ) (373) = 30.6 L for one mol of water vapor P 1.00 Since the density of H2O(l) is 1.00 g/cm3, 1.00 mol of H2O(l) occupies 18.0 cm3 or 0.0180 L. w = 1.00 atm (30.6 L 0.0180 L) = 30.6 L atm w = 30.6 L atm × 101.3 J/Latm = 3.10 × 103 J = 3.10 kJ 368 CHAPTER 10 LIQUIDS AND SOLIDS ΔE = q + w = 40.7 kJ 3.10 kJ = 37.6 kJ 37.6 × 100 = 92.4% of the energy goes to increase the internal energy of the water. 40.7 The remainder of the energy (7.6%) goes to do work against the atmosphere. 119. For a cube: (body diagonal)2 = (face diagonal)2 + (cube edge length)2 In a simple cubic structure, the atoms touch on cube edge so the cube edge = 2 r where r = radius of sphere. 2r 2r face diagonal = body diagonal = (2 r ) 2 (2 r ) 2 = 4 r2 4 r2 = r 8 = 2 2 r (2 2 r ) 2 (2 r ) 2 = 12 r 2 = 2 3 r The diameter of the hole = body diagonal 2 radius of atoms at corners. diameter = 2 3 r 2 r; Thus, the radius of the hole is: 4 2 3 2 The volume of the hole is: π r 3 2 120. 2 3 r 2 r 2 3 2 r 2 2 3 C2H6O, 2(4) + 6(1) + 6 = 20 e; Determining the two possible structures is trial and error. Because hydrogen is always attached with a single bond, H cannot be a central atom. The two possible structures having two bonds and two lone pairs for O are: H H H C C H H H O liquid (H-bonding) H and H C H H O C H gas (no H-bonding) H CHAPTER 10 LIQUIDS AND SOLIDS 369 The first structure with the –OH bond is capable of forming hydrogen bonding; the other structure is not. Therefore, the liquid (which has the stronger intermolecular forces) is the first structure and the gas is the second structure. Integrative Problems 121. 19.0 g = 144 g/mol 0.132 mol molar mass XY = X: [Kr] 5s24d10; This is cadmium, Cd. molar mass Y = 144 – 112.4 = 32 g/mol; Y is sulfur, S. The semiconductor is CdS. The dopant has the electron configuration of bromine, Br. Because Br has one more valence electron than S, doping with Br will produce an n-type semiconductor. 122. Assuming 100.00 g of MO2: 23.72 g O × 1 mol O = 1.483 mol O 16.00 g O 1.483 mol O × 1 mol M = 0.7415 mol M 2 mol O 100.00 g – 23.72 g = 76.28 g M molar mass M = 76.28 g = 102.9 g/mol 0.7415 mol From the periodic table element M is rhodium, Rh. The unit cell for cubic closest packing is face-centered cubic (4 atoms/unit cell). The atoms for fcc are assumed to touch along the face diagonal of the cube so face diagonal = 4 r. The distance between the centers of touching Rh atoms will be the distance of 2 r where r = radius of Rh atom. face diagonal = 2 l, where l = cube edge face diagonal = 4 r = 2 × 269.0 × 10 12 m = 5.380 × 10 10 m 2 l = 4 r = 5.38 × 10 10 m, l = 4 atoms Rh density = 5.38 10 10 m 2 = 3.804 × 10 10 m = 3.804 × 10 8 cm 1 mol Rh 102 .9 g Rh 23 mol Rh 6.0221 10 atoms = 12.42 g/cm3 8 3 (3.804 10 cm) 370 123. CHAPTER 10 P ln 1 P2 LIQUIDS AND SOLIDS ΔH vap 1 1 296 J 200 .6 g ; ΔH vap = 5.94 × 104 J/mol Hg R g mol T T 1 2 2.56 10 3 torr 5.94 10 4 J / mol ln 8.3145 10 4 J / K mol P2 1 1 573 K 298 .2 K 2.56 10 3 torr = 11.5, P2 = 2.56 × 10 3 torr / e 11.5 = 253 torr ln P 2 1 atm 253 torr 15.0 L 760 torr PV n= = 0.106 mol Hg 0.08206 L atm RT 573 K K mol 0.106 mol Hg × 6.022 10 23 atoms Hg = 6.38 × 1022 atoms Hg mol Hg Marathon Problem 124. q = s × m × T; heat loss by metal = heat gain by calorimeter. The change in temperature for the calorimeter is: T = 25.2C 0.2C – 25.0C 0.2C. Including the error limits, T can range from 0.0C to 0.4C. Because the temperature change can be 0.0C, there is no way that the calculated heat capacity has any meaning. The density experiment is also not conclusive. d= 4g = 10 g/cm3 (1 significant figure) 3 0.42 cm dhigh = 5g = 12.5 g/cm3 = 10 g/cm3 to 1 sig fig 3 0.40 cm dlow = 3g = 7 g/cm3 3 0.44 cm From Table 1.5, the density of copper is 8.96 g/cm3. The results from this experiment cannot be used to distinguish between a density of 8.96 g/cm3 and 9.2 g/cm3. The crystal structure determination is more conclusive. Assuming the metal is copper: 3 1 10 10 cm 2.16 10 22 cm3 volume of unit cell = (600. pm) 1 pm 3 CHAPTER 10 LIQUIDS AND SOLIDS Cu mass in unit cell = 4 atoms d= 371 1 mol Cu 63.55 g Cu = 4.221 × 10 22 g Cu 23 mol Cu 6.022 10 atoms mass 4.221 10 22 g = 1.95 g/cm3 volume 2.16 10 22 cm3 Because the density of Cu is 8.96 g/cm3, then one can assume this metal is not copper. If the metal is not Cu, then it must be kryptonite (as the question reads). Because we don’t know the molar mass of kryptonite, we cannot confirm that the calculated density would be close to 9.2 g/cm3. To improve the heat capacity experiment, a more precise balance is a must and a more precise temperature reading is needed. Also, a larger piece of the metal should be used so that T of the calorimeter has more significant figures. For the density experiment, we would need a more precise balance and a more precise way to determine the volume. Again, a larger piece of metal would help in order to insure more significant figures in the volume.