Chapter 10 Liquids and Solids Some unusual properties of water, such as its high boiling point and high molar heat capacity, allow it to exist as a liquid on Earth, to moderate the climate, and to support life. The theme of this chapter is the effect of the nature of molecules (or the individual atoms or ions) on the structure and properties of the liquids and solids they form.

Figure 10.1 London forces arise from the attraction between two instantaneous dipoles. The dipoles are due to fluctuations in electron locations in the molecules. Although the instantaneous dipoles on two molecules are continuously changing direction, they remain in step long enough to attract each other.

Figure 10.2 Hydrocarbons show how London forces increase in strength with molar mass. Pentane, C 5 H 12, is a mobile fluid (left); pentadecane, C 15 H 32, a viscous liquid (middle); and octadecane, C 18 H 38, a waxy solid (right). The effect of increasing intermolecular forces is enhanced by the ability of long-chain molecules to become entangled with one another.

Figure 10.3 (a) The instantaneous dipole moments (represented by the color gradient) in two neighboring rod-shaped molecules tend to be close together and to interact strongly. (b) Those on neighboring spherical molecules tend to be far apart and to interact weakly.

Figure 10.4 Polar molecules attract each other by the interaction between the permanent partial charges of their electric dipoles (represented by the arrows). Both the relative orientations shown (end-to-end and side-by-side) are energetically favorable.

Figure 10.5 The boiling points of the hydrogen compounds of most of the p-block elements show a smooth variation and increase with molar mass in each group. However, three compounds—ammonia, water, and hydrogen fluoride—are strikingly out of line. This anomaly suggests the existence of strong intermolecular interactions not present in the other compounds shown.

Figure 10.6 The relative viscosities of several liquids, compared with water. Liquids composed of molecules that cannot form hydrogen bonds are generally less viscous than those that can form hydrogen bonds. Mercury is an exception: its atoms stick together by a kind of metallic bonding, and its viscosity is relatively high.

Figure 10.7 Surface tension arises from the attractive forces acting on the molecules at the surface, as shown in the inset. A molecule within the liquid experiences forces from all directions, but a molecule at the surface experiences a net inward force.

Figure 10.8 The nearly spherical shape of these beads of water on the waxy surface of a leaf arises from the effect of surface tension. The beads are flattened slightly by the effect of the Earth’s gravity.

Figure 10.9 When the adhesive forces between a liquid and glass are stronger than the cohesive forces within the liquid, the liquid forms a concave meniscus, as shown here for water in glass (left). When the cohesive forces are stronger than the adhesive forces (as they are for mercury in glass), the surface is convex, curved downward (right).

Figure Crystalline solids have well-defined faces and an orderly internal structure. Each face is the edge of a stack of atoms, molecules, or ions. The crystal in the photograph is galena, PbS.

Figure (a) Quartz is a crystalline form of silica, SiO 2 ; and the atoms are arranged in an orderly network, represented in two dimensions in the overlay. (b) When molten silica solidifies rapidly, it becomes glass. Now the atoms form a disorderly network.

Investigating Matter 10.1 (A)(a) Constructive interference between two waves (the red and green lines), which gives a wave of greater amplitude (the orange line); (b) destructive interference between the same two waves, to give a wave of smaller amplitude.

Investigating Matter 10.1 (B) The diffraction pattern formed when x- rays pass through a single crystal of sodium chloride. The bright spots are the photographic record of the pattern formed where x-rays interfere constructively.

Investigating Matter 10.1 (C) The distance d between layers of atoms determines the angle  at which the x-rays reflected from the two layers are in phase and interfere constructively.

Case Study 10 (a): A scanning tunneling microscope image of liquid crystal molecules aligned in a smectic phase.

Case Study 10 (b) : (a) In the nematic phase of a liquid crystal, the long molecules lie parallel to one another. (b) In the smectic phase of a liquid crystal, not only do the molecules lie parallel to one another, they also form sheets. (c) In the cholesteric phase of a liquid crystal, layers of parallel molecules are rotated relative to their neighbors and form a helical structure.

Case Study 10 (c): A cross section through one of the two layers of a cell membrane. The long, narrow molecules are aligned with their polar heads toward the surfaces of the membrane.

Figure A close-packed structure can be built up in stages. The first layer (A) is laid down with minimum waste of space, and the second layer (B) lies in one-half the dips—the depressions—of the first. The inset shows a top-down view.

Figure The third layer of spheres can be deposited so that the atoms lie directly above the atoms of the first layer to give an ABABAB... structure. Notice in the top-down view (inset) that the holes in each layer are lined up so that in certain places (indicated by the arrow) you can see all the way through the crystal. You can also see this effect when you stack pennies in an ABAB structure.

Figure A fragment of the structure formed as described in Fig shows the hexagonal symmetry of the arrangement, and the origin of its name “hexagonal close-packed.”

Figure As an alternative to the scheme shown in Fig , the spheres of the third layer can lie in the dips above the dips in the first layer to give an ABCABC... arrangement of layers. In this case, you cannot see through the layers because each hole has an atom either on top of it or below it.

Figure A fragment of the structure formed as described in Fig shows the origin of the names “cubic close-packed” or “face- centered cubic” for this arrangement. The layers A, B, C can be seen along the diagonals and are indicated by the different colors of atoms.

Figure The body-centered cubic (bcc) structure. This structure is not packed as closely as the others we have illustrated. It is less common among metals than the close-packed structures, but some ionic structures are based on it.

Figure A primitive cubic structure (which later we see is a unit cell) has an atom at each of the eight corners.

Figure The entire crystal structure is constructed from a unit cell by stacking the cells together without any gaps in between. In this cubic close-packed (face-centered cubic) cell, each corner atom is shared by eight cubes that touch at the corner. Each atom on a face is shared by two adjacent cubes.

Figure The unit cells of the (a) ccp (or fcc) and (b) bcc structures in which the locations of the centers of the spheres are marked by dots.

Figure The contribution 1/8 or 1/2 of the individual atoms in a face-centered cell to the total number of atoms in a unit cell.

Figure The geometries of three cubic unit cells, showing the relation of the dimensions of each cell to the radius of a sphere representing an atom or ion, r. The side of a cell is a; the diagonal of the body of a cell, b; and the diagonal of a face, f. (a) Primitive cubic cell; (b) body-centered cubic cell; (c) face-centered cubic cell.

Figure (a) When light of a particular color shines on the surface of a metal, the electrons at the surface oscillate in step. This oscillating motion gives rise to an electromagnetic wave that we perceive as the reflection of the source.

Figure (b) Each of these solar mirrors in California is positioned at the best angle to reflect sunlight of all wave-lengths into a collector that uses it to generate electricity.

Figure (a) A metal is malleable because, when cations are displaced by a blow from a hammer, the mobile electrons can immediately respond and follow the cations to their new positions. (b) This piece of lead has been flattened by hammering, whereas crystals of ionic compounds, such as the orange compound lead(II) oxide shown here, shatter when struck.

Figure A metallic conductor is a substance with a resistance that increases with temperature. A semiconductor is a substance with a resistance that decreases with increasing temperature. A superconductor is a substance that has zero resistance below a certain temperature. An insulator behaves like a semiconductor with a very high resistance.

Figure A metallic solid contains so many bonded atoms that their molecular orbitals form an almost continuous band. At the lower edge of the band, the molecular orbital is fully bonding; at the upper edge, the molecular orbital is fully antibonding. The inset shows that, although the band of allowed energies appears to be continuous, it is in fact composed of discrete, closely spaced levels.

Figure The band structure typical of a variety of solids. The occupied regions of the bands are shaded tan. Note that the distinction between a solid insulator and a semiconductor is the width of the gap between the valence and conduction bands.

Figure This solar-powered vehicle makes use of silicon-based semiconductors to convert sunlight into electrical energy.

Figure In a substitutional alloy, the positions of some of the atoms of one metal are taken by atoms of another metal. The two elements must have similar atomic radii.

Figure In an interstitial alloy, the atoms of one metal lie in the gaps between the atoms of another metal. The two elements must have markedly different atomic radii.

Figure Many racing bicycle frames like this one are built from high-strength, low-density steels formed by alloying iron with metals such as manganese, molybdenum, and titanium.

Figure The arrangement of ions in the rock-salt structure. (a) The unit cell, showing the packing of the individual ions. (b) A representation of the structure in terms of dots that identify the centers of the ions. The pink spheres are cations and the green spheres are anions.

Figure The smooth faces of the crystals of sodium chloride shown in this micrograph result when billions of unit cells stack together in an orderly way. The first inset (top) shows some of the stacked unit cells, viewed from one side of the crystal. The middle inset identifies the individual ions. The third inset (lower right) illustrates the coordination of an anion to its six cation neighbors.

Figure The cesium-chloride structure: (a) the unit cell and (b) the locations of the centers of the ions. The pink spheres are cations and the green spheres are anions.

Figure The repetition of the cesium-chloride unit cell recreates the entire crystal. This view from one side of the crystal shows six unit cells.

Figure The structure of nickel(II) arsenide. Structures like this are found when covalent bonding is starting to become important and the ions have to take up specific positions relative to one another to maximize their bonding.

Figure As the crystalline structure of ice forms, the regular array of molecules creates intricate designs. The water molecules in ice are held together by hydrogen bonds in a relatively open structure. Each O atom (red) is surrounded tetrahedrally by four hydrogen atoms, two of which are s-bonded to it and two of which are hydrogen bonded to it.

Figure As a result of an open structure, ice is less dense than water and floats in it (left). Benzene molecules can pack more tightly than water molecules in the solid state and, as a result, solid benzene is denser than liquid benzene; “benzenebergs” sink in liquid benzene (right).

Figure Part of the structure of a diamond. Each sphere represents the location of the center of a carbon atom. Each atom forms an sp 3 hybrid covalent bond to each of its four neighbors. The highly regular structure results in the smooth crystal faces seen in the photograph.

Figure Graphite consists of staggered layers of hexagonal rings of sp 2 hybridized carbon atoms. On the left is a view along the layers; on the right, a view perpendicular to them. The slipperiness of graphite is due to the ease with which the layers can slide over one another when impurity atoms are present.

Figure The apparatus on the left shows a mercury barometer, with a vacuum over the mercury column and a flask containing a liquid and its vapor. The stopcock between the top of the barometer and flask is closed. In the apparatus on the right, the stopcock has been opened and the vapor pressure exerted by the liquid in the flask (in Torr) can be measured by recording the distance in millimeters by which the column of mercury has been lowered. The vapor pressure is the same, however much liquid is present.

Figure Three tubes of mercury; the tube in the middle is attached to a flask of water like that in Fig , and the one on the right is similarly attached to a flask of ethanol. Neither flask is shown. The fact that ethanol has a higher vapor pressure than water is shown by the fact that the level of mercury in the tube connected to the ethanol flask is lower than the level in the tube connected to the flask with water. The total pressure at the base of each column is the sum of the pressure due to the weight of the mercury and the vapor pressure of the liquid.

Figure When a liquid and its vapor are in dynamic equilibrium inside a closed container, the rate at which the molecules leave the liquid is equal to the rate at which they return.

Figure The variation of the vapor pressure of liquids with temperature, for diethyl ether (orange), ethanol (red), benzene (green), and water (blue). The normal boiling point is the temperature at which the vapor pressure is 1 atm (760 Torr).

Figure Phase diagram for water (not to scale). The solid lines define the boundaries of the regions of pressure and temperature at which each phase is the most stable. Note that the freezing point decreases with increasing pressure. The letters A and B are referred to in Example 10.7.

Figure The changes undergone by a liquid as its pressure is decreased at constant temperature. The dot on the vertical line traces the path taken by the system, which is described in the text.

Figure Phase diagram for carbon dioxide. Note that carbon dioxide sublimes to a vapor at 1 atm. Liquid carbon dioxide can exist only at pressures above 5.1 atm.

Figure Phase diagram for sulfur. Notice that there are three triple points and that the pressure scale covers a very wide range of values.

Figure The phase diagram for a one-component system and the interpretation of each region, line, and point.

Figure The phase diagram for water shown with more detail. The triple point and critical point are labeled.

Figure When the temperature of a liquid in a sealed, constant- volume container is raised, the density of the liquid decreases and the density of the vapor increases as the liquid evaporates. At the critical temperature, T c, the density of the vapor becomes the same as the density of the liquid; above that temperature, a single uniform phase fills the container.