Solids intermolecular forces

The most important manifestation of intermolecular forces is the very existence of liquids and solids. Without intermolecular forces, solids and liquids would not exist and all matter would be gaseous. In liquids, we can observe several other manifestations of intermolecular forces including surface tension and viscosity. 

As also noted in the preceding chapter, it is customary to divide adsorption into two broad classes, namely, physical adsorption and chemisorption. Physical adsorption equilibrium is very rapid in attainment (except when limited by mass transport rates in the gas phase or within a porous adsorbent) and is reversible, the adsorbate being removable without change by lowering the pressure (there may be hysteresis in the case of a porous solid). It is supposed that this type of adsorption occurs as a result of the same type of relatively nonspecific intermolecular forces that are responsible for the condensation of a vapor to a liquid, and in physical adsorption the heat of adsorption should be in the range of heats of condensation. Physical adsorption is usually important only for gases below their critical temperature, that is, for vapors. 

Surface Protection. The surface properties of fluorosihcones have been studied over a number of years. The CF group has the lowest known intermolecular force of polymer substituents. A study (6) of liquid and solid forms of fluorosihcones has included a comparison to fluorocarbon polymers. The low surface tensions for poly(3,3,3-trifluoropropyl)methylsiloxane and poly(3,3,4,4,5,5,6,6,6-nonafluorohexyl)methylsiloxane both resemble some of the lowest tensions for fluorocarbon polymers, eg, polytetrafluoroethylene. 

Liquid-solid interactions due to long-range intermolecular forces are much larger than are gas-solid interactions. This means that it is easier to collect fine particles at a liquid-liquid interface than at a gas-liquid interface. 

It follows from this discussion that all of the transport properties can be derived in principle from the simple kinetic dreoty of gases, and their interrelationship tlu ough k and c leads one to expect that they are all characterized by a relatively small temperature coefficient. The simple theory suggests tlrat this should be a dependence on 7 /, but because of intermolecular forces, the experimental results usually indicate a larger temperature dependence even up to for the case of molecular inter-diffusion. The Anhenius equation which would involve an enthalpy of activation does not apply because no activated state is involved in the transport processes. If, however, the temperature dependence of these processes is fitted to such an expression as an algebraic approximation, tlren an activation enthalpy of a few kilojoules is observed. It will thus be found that when tire kinetics of a gas-solid or liquid reaction depends upon the transport properties of the gas phase, the apparent activation entlralpy will be a few kilojoules only (less than 50 kJ). 

A useful property of liquids is their ability to dissolve gases, other liquids and solids. The solutions produced may be end-products, e.g. carbonated drinks, paints, disinfectants or the process itself may serve a useful function, e.g. pickling of metals, removal of pollutant gas from air by absorption (Chapter 17), leaching of a constituent from bulk solid. Clearly a solution s properties can differ significantly from the individual constituents. Solvents are covalent compounds in which molecules are much closer together than in a gas and the intermolecular forces are therefore relatively strong. When the molecules of a covalent solute are physically and chemically similar to those of a liquid solvent the intermolecular forces of each are the same and the solute and solvent will usually mix readily with each other. The quantity of solute in solvent is often expressed as a concentration, e.g. in grams/litre. 

Van der Waals forces (Section 2.13) Intermolecular forces that are responsible for holding molecules together in the liquid and solid states. 

When iodine chloride is heated to 27°C, the weak intermolecular forces are unable to keep the molecules rigidly aligned, and the solid melts. Dipole forces are still important in the liquid state, because the polar molecules remain close to one another. Only in the gas, where the molecules are far apart, do the effects of dipole forces become negligible. Hence boiling points as well as melting points of polar compounds such as Id are somewhat higher than those of nonpolar substances of comparable molar mass. This effect is shown in Table 9.3. 

We have seen that the pure elements may solidify in the form of molecular solids, network solids, or metals. Compounds also may condense to molecular solids, network solids, or metallic solids. In addition, there is a new effect that does not occur with the pure elements. In a pure element the ionization energies of all atoms are identical and electrons are shared equally. In compounds, where the most stable electron distribution need not involve equal sharing, electric dipoles may result. Since two bonded atoms may have different ionization energies, the electrons may spend more time near one of the positive nuclei than near the other. This charge separation may give rise to strong intermolecular forces of a type not found in the pure elements. 

A common feature of all clathrates discussed so far is a host lattice, by itself thermodynamically unstable, which is stabilized by inclusion of the second component. The forces binding this component must be similar in nature to the intermolecular forces in liquids. It seems natural, therefore, to regard a clathrate compound as a solid solution of the second component in the (meta-stable) host lattice. 

In Chapter 4 we considered gases, in which intermolecular forces play only a minor role. Here, we deal with liquids and solids, in which the forces that hold molecules together are of crucial importance for determining the physical properties of bulk samples. Individual water molecules, for instance, are not wet, but bulk water is wet because water molecules are attracted to other substances and spread over their surfaces. Individual water molecules neither freeze nor boil, but bulk water does, because in the process of freezing molecules stick together and form a rigid array and in boiling they separate from one another and form a gas. 

We have to refine our atomic and molecular model of matter to see how bulk properties can be interpreted in terms of the properties of individual molecules, such as their size, shape, and polarity. We begin by exploring intermolecular forces, the forces between molecules, as distinct from the forces responsible for the formation of chemical bonds between atoms. Then we consider how intermolecular forces determine the physical properties of liquids and the structures and physical properties of solids. 

Intermolecular forces are responsible for the existence of several different phases of matter. A phase is a form of matter that is uniform throughout in both chemical composition and physical state. The phases of matter include the three common physical states, solid, liquid, and gas (or vapor), introduced in Section A. Many substances have more than one solid phase, with different arrangements of their atoms or molecules. For instance, carbon has several solid phases one is the hard, brilliantly transparent diamond we value and treasure and another is the soft, slippery, black graphite we use in common pencil lead. A condensed phase means simply a solid or liquid phase. The temperature at which a gas condenses to a liquid or a solid depends on the strength of the attractive forces between its molecules. 

The liquid phase of matter is the most difficult to visualize. We have seen that a gas-phase molecule moves with almost complete freedom. The intermolecular forces from other molecules are minimal, and movement is highly disordered. In the solid phase, a molecule is locked in place by intermolecular forces and can only oscillate around an average location. The liquid phase lies between the extremes of the gas and solid phases. The molecules are mobile, blit they cannot escape from one another completely. 

Molecular solids are assemblies of discrete molecules held in place by intermolecular forces. 

Account for the following observations in terms of the type and strength of intermolecular forces, (a) The melting point of solid xenon is —112°C and that of solid argon is — 189°C. 

Crystal engineering. Utilization of noncovalent intermolecular forces in the solid state to design new nanomaterials with desired functions. 

Every gas changes into a liquid if the pressure is high enough and the temperature is low enough. The atoms or molecules of a liquid or solid stick together in a finite volume rather than expanding, as a gas does, to fill all available space. This cohesiveness comes from electrical forces of attraction between the negative electron cloud of each atom and the positive nuclei of other atoms. We describe intermolecular forces in Chapter 11. 

Forces of attraction between molecules are responsible for the existence of liquids and solids. In the absence of these intermolecular forces, all molecules would move independently, and all substances would be gases. The natural phases of the elements indicate the importance of intermolecular forces. At room temperature and pressure, only 11 elements are gases. Mercuiy and bromine are liquids, and all the rest of the elements are solids. For all but the 11 gaseous elements, intermolecular forces are too large to ignore under normal conditions. 

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